II 






II 



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COPYRIGHT DEPOSIT. 



FIRST YEAR CHEMISTRY 



A TEXT IN ELEMENTARY CHEMISTRY 
FOR SECONDARY SCHOOLS 



BY 

WILHELM SEGERBLOM, A.B. 

INSTRUCTOR IN CHEMISTRY 
AT THE PHILLIPS EXETER ACADEMY 



N.l/ 



PROCFiESSVS 
REVELAT.IOKE 



EXETER, N. H. 
EXETER BOOK PUBLISHING COMPANY 

1909 



Q>, 



& 
'<*>* 



Copyright, 1909 
By Wilhelm Segerblom 



:' 8 4 



The News-Letter Press 
Exeter, N. H. 



PREFACE 



This book is not a synopsis of the whole of chemistry; 
nor is it a collection of sample fragments from the various 
departments of this rapidly widening subject. It em- 
braces as much of the essential facts and influences of chemis- 
try as the ordinary student can assimilate in one year's 
study of the subject, presented in the order that the human 
mind — particularly the student mind — exploring new re- 
gions naturally works, and tied into a logical and connected 
whole. 

Briefly, my plan is: To use the inductive method rigor- 
ously in the first ninety experiments — about one-third of 
the book — beginning with the metals instead of with the 
nonmetals, in the belief that the student can get more tangi- 
ble results from the more tangible substances. The last 
four of these ninety experiments contain a short chemical 
investigation of a substance whose composition is unknown 
to the student at this stage of the year's work. The per- 
forming of these ninety experiments reveals to the student 
the methods of thought used by the chemical investigator 
in attacking his problems, — the same method that he (the 
student) must use on a small scale in solving his own diffi- 
culties; the result of all this is that the student is taught 
to think his own way out of all his difficulties. Here I 
drop the strictly inductive method, for I realize that it can- 



iv PREFACE 

not be used to advantage in teaching theory and the large 
amount of descriptive matter that comes later. The in- 
vestigation is followed by the theory of chemistry, with 
experiments to illustrate the laws and principles; this 
comes at the middle of the } r ear, — a period now approved 
by many teachers. The theory of chemistry is followed 
by considerable descriptive chemistry, studied in the light 
of present theoretical conceptions, but with the spirit of 
the inductive method still an unconscious guide. 

The author realizes that he has struck out into unbeaten 
paths and may, perhaps, be going in cross directions to the 
trend of many chemistry teachers, but he feels confident 
of the correctness of his mental compass, for he has watched 
carefully the trial trips of the text in his own classes during 
the past four years; and, after all, the worth and value of a 
text lies in its ability to lead the confident and confiding 
student mind naturally and progressively up towards the 
hights of vision already attained by the teacher. 

My object has been to give the student a conception of 
chemistry from the point of view of the scholar and the 
thinker, rather than from the point of view of the crammer, 
i.e., to teach chemistry rather than to teach a text-book. 
It is generally accepted that the chemistry of the labora- 
tory is the only real chemistry, and that the atom and the 
ion are simply helps in discussing facts; and for this pres- 
entation of the subject the deductive method is as essential 
as the inductive. 

As lie Bon says: "No physical law is now capable of 
exact verification. All are merely approximations to the 
truth, stating what would be the case if all disturbing fac- 
tors were removed." Be that as it may, every thoughtful 
teacher of laboratory science knows how difficult it is to 
eliminate these " disturbing factors" from the experiments 
that are to be presented to the student, and then after the 



PREFACE V 

elimination to have anything left worth presenting that 
can be worked into a united whole. The author will doubt- 
less be pardoned for saying that he first laid out the logical 
plan of experimental chemistry presented in these pages, 
but that by far the bulk of his work has been to adapt the 
work to the ability and limitations of the beginner by 
eliminating and circumventing as many as possible of those 
''disturbing factors" that lead to secondary and confusing 
phenomena. 

The paragraphs in fine print at the end of many experi- 
ments answer briefly just those questions which arise in 
the alert student mind but which cannot easily be answered 
from the student's personal observation in the laboratory. 
Each teacher must of course use his own judgment in the 
use of the fine print notes, elaborating them, if he thinks 
it necessary, to the extent that each class is ready for and 
seeks such elaboration. The author finds it advisable not 
to go into great detail in such elaboration, — at least not 
in the first ninety experiments. 

William A. Noyes, when speaking of the contribution of 
chemistry to modern life in his recent inaugural address 
at the University of Illinois, said: " While the material 
advantages which have come to us from chemistry are 
very great and may be justly emphasized, its greatest 
achievement is, after all, the part which it has had, to- 
gether with other sciences, in transforming the way in 
which the world thinks. In its laboratory method it has 
replaced the old idea of authority by the idea of first-hand 
knowledge. It leads the individual to seek for himself 
the fundamental basis of his knowledge and it leads him 
not merely to pass knowledge on to the next generation, 
but to transform it into a new and truer form. And as 
this scientific spirit permeates society, it more and more 
destroys deceit and fraud, wherever found." If the author 



vi PREFACE 

has succeeded in furthering this spirit of improved think- 
ing, then his labors in preparing this little text have not 
been in vain. 

Thanks are gratefully accorded those of my pupils and 
other friends who by criticism or suggestion have aided in 
perfecting these pages. Mr. W. A. Stone, of Albany, N. Y., 
has read the manuscript; Mr. J. B. Merrill, of Woonsocket, 
R. I., has read both manuscript and proof; their kindly 
interest and helpful suggestions are gladly acknowledged. 
To my former teacher, Dr. George R. White, of Wellesley 
Hills, Mass., I am most deeply indebted; during my student 
days his sympathetic teaching opened up to me the allur- 
ing field of chemistry, and without that introduction the 
development of this little text and the scientific hopefulness 
which has made its preparation more a pleasure than a 
burden would not have been possible. The author assumes 
the responsibility for the statements in this book. Any cor- 
rections or comments on the text will be gratefully received. 

W. S. 

Exeter, N. H., 1909. 



CONTENTS 



(The references are to pages.) 

Introduction 1 

Apparatus ' . 1 

Laboratory Notebook 1 

Preliminary Work 3 

Experiment 1. Measuring 4 

Metric System 4 

Measurement of desk 5 

Measurement of glass vessels 6 

How to read the graduate 7 

Purpose of this work in measuring 8 

Experiment 2. Weighing 8 

Platform balance 8 

Set of iron weights . 9 

How to weigh 9 

Horn-pan balance 10 

Set of smaller weights 10 

Heavy weighing 11 

Light weighing 11 

Tare weighing 12 

Experiment 3. Manipulating glass 12 

Bat-wing burner 12 

Bunsen burner 13 

How to cut glass tubing 14 

How to fire-polish glass tubing 15 

How to bend glass tubing 15 

How to draw out glass tubing 16 

Experiment 4. Making a wash-bottle 17 

Experiment 5. Copper and its properties 19 

Definition of property 19 

List of properties . . . 20 

How to test for properties 20 

The blast lamp and its use 23 

How to light the blast lamp 24 

Experiment 6. Heating copper in contact with air . . 25 

vii 



viii CONTENTS 

(The references are to pages.) 
Experiment 6 — Continued 

Oxidation 27 

Note on copper 27 

Experiment 7. Zinc and its properties 27 

Experiment 8. Heating zinc in contact with air ... 27 

Note on zinc 28 

Experiment 9. Magnesium and its properties .... 28 

Experiment 10. Heating magnesium in contact with air . . 28 

Note on magnesium 30 

Experiment 11. Phosphorus and its properties .... 30 

Caution 30 

Melting point of phosphorus 31 

Experiment 12. Heating phosphorus in contact with air . . 32 

Proportion by volume of oxygen in air 34 

General Considerations 35 

Definition of simple substance 35 

Definition of compound substance 36 

Physical changes and chemical changes 36 

Equations 37 

Factors and products . . 38 

Note on phosphorus 38 

Air and its composition . . 39 

Experiment 13. Tron and its properties 40 

Experiment 14. Heating iron in contact with air 40 

Note on iron 41 

Experiment 15. Mercury and its properties .... 41 

Surface tension 42 

Experiment 16. Mercury oxide 43 

Experiment 17. Preparation of oxygen from mercury oxide . 43 

How to make a bulb tube . . . . . . . 43 

Definition of analysis 45 

Note on mercury 45 

Experiment 18. Preparation of oxygen from potassium chlorate 45 
Experiment 19. Preparation of oxygen from potassium chlorate 

on a large scale 46 

How to filter . 49 

Definition of filtrate 51 

Definition of precipitate 51 

Washing precipitates 52 

Drying precipitates ,.,.,., 52 



CONTENTS ix 

(The references are to pages.) 

Experiment 19 — Continued 

Heating to constant weight 53 

Purpose of the black oxide of manganese .... 53 

Catalytic action 53 

Catalytic agent 53 

Experiment 20. Action of undiluted oxygen gas on zinc . . 54 

Experiment 21. Action of undiluted oxygen gas on magnesium 54 

Experiment 22. Action of undiluted oxygen gas on phosphorus 55 

Note on oxygen 55 

Ozone 55 

Oxidation and kindred topics ....... 55 

Combustion 56 

Kindling temperature 56 

Spontaneous combustion 56 

Experiment 23. The decomposition of water by means of elec- 
trolysis 56 

Experiment 24. The volumetric composition of water . . 60 

Note on the RhumkorrT coil 65 

Experiment 25. Preparation of hydrogen from water by means 

of magnesium 65 

Definition of reduction ,68 

Reducing agent 68 

Note on water 69 

Experiment 26. Sulphur and its properties .... 70 

Experiment 27. Modifications of sulphur ..... 70 

How to use the magnifying glass 71 

Note on sulphur 73 

Allotropy 73 

Experiment 28. Heating sulphur in contact with air and in con- 
tact with oxygen 74 

Catching gases by displacement of air 74 

Experiment 29. Preparation of sulphurous acid ... 76 

Component parts ,77 

Note on litmus paper 77 

Note on sulphurous acid 78 

Experiment 30. Heating sulphur oxide in contact with oxygen 78 

Catch bottles 78 

Experiment 31. Preparation of sulphuric acid .... 83 

Note on sulphuric acid 84 

The lead chamber process 84 



x , CONTENTS 

(The references are to pages.) 

Experiment 31 — Continued 

The contact process 85 

Experiment 32. Preparation of hydrogen from sulphuric acid, 

or the action of sulphuric acid on zinc .... 85 

Making a generator 85 

Generating hydrogen 87 

Sealing the thistle tube 88 

Lighting hydrogen with the safety tube 88 

Note on the color of the hydrogen flame 90 

Treatment of the residue in the flask 90 

Crystallization 91 

Composition of the crystals 91 

Water of crystallization 93 

Efflorescence 94 

Crystallized zinc sulphate 94 

Anhydrous zinc sulphate 94 

Note on zinc sulphate 95 

Note on sulphates in general 95 

Experiment 33. Action of sulphuric acid on iron ... 95 

Generating hydrogen 95 

Treatment of the residue in the flask 96 

Composition of the crystals . . ■ 97 

Definition of substitution 99 

Note on iron sulphate 99 

Experiment 34. Action of sulphuric acid on magnesium . . 99 

Generating hydrogen 99 

Treatment of the residue in the flask 100 

Composition of the crystals 100 

Note on magnesium sulphate 101 

Definition of the term salt 101 

Experiment 35. Action of sulphuric acid on copper . . . 101 

Generating sulphur dioxide 101 

Treatment of the residue in the flask 102 

Definition of decantation 102 

Composition of the crystals . 103 

Recrystallization . 103 

Note on copper sulphate 104 

Experiment 36. The action of sulphuric acid on zinc oxide . 104 

Note on the action of sulphuric acid on oxides . . . 105 

Experiment 37. Reaction between copper and sulphur . . 105 



CONTENTS xi 

(The references are lo pages.) 

Experiment 3S. Reaction between mercury and sulphur . . 107 

Definition of trituration 107 

Experiment 39. Reaction between zinc and sulphur . . . 107 

Experiment 40. Reaction between iron and sulphur . . . 108 

Definition of synthesis 109 

Experiment 41. Reaction between hydrogen and warm sulphur 110 

Experiment 42. Action of sulphuric acid on iron sulphide . 112 

Experiment 43. Formation of a metallic sulphide by precipita- 
tion 113 

Experiment 44. Action of sulphuric acid on zinc sulphide . 114 
Note on the action of sulphuric acid on sulphides . . .115 

Experiment 45. Carbon, its properties, and its allotropic forms 115 

Note on carbon 115 

Experiment 46. Heating carbon in contact with oxygen . .116 
Study of carbon oxide and of carbonic acid . . . .118 

Experiment 47. Action of magnesium on carbon oxide .119 

Experiment 48. Action of hot zinc on carbon oxide . . . 120 

Composition of the new gas 122 

Carbon monoxide 123 

Carbon dioxide 123 

Note on the oxides of carbon 123 

Definition of the term anhydride 123 

Note on carbonic acid 123 

Modification of the work on the oxides of carbon . . . 124 

Experiment 49. Chlorine and its properties .... 124 

Note on chlorine 125 

Experiment 50. Reaction between hydrogen and chlorine . 126 

Note on hydrogen chloride 127 

Experiment 51. Action of hydrochloric acid on zinc . . 127 

Making a test tube generator 127 

Experiment 52. Action of hydrochloric acid on iron . . 128 

Experiment 53. Action of hydrochloric acid on magnesium . 129 

Experiment 54. Sodium and its properties 129 

Note on sodium 130 

Experiment 55. Heating sodium in contact with air . . 130 

Sodium monoxide 131 

Sodium dioxide 131 

Experiment 56. Action of water on sodium oxide . . . 131 

Definition of alkaline 132 

Definition of deliquescence 132 



xii CONTENTS 

(The references are to pages.) 
Experiment 56 — Continued 

Note on sodium hydroxide 133 

Note on turmeric paper 133 

Experiment 57. The action of sodium on water . . . 134 

Twin equations 135 

Experiment 58. Action of chlorine on sodium .... 136 
Experiment 59. Reaction between sodium hydroxide and 

sulphuric acid 137 

Neutralization 138 

Note on sodium sulphate 139 

Experiment 60. Reaction between sodium hydroxide and hy- 
drochloric acid 139 

Note on sodium chloride . 140 

Experiment 61. Reaction between sodium hydroxide and car- 
bonic acid . . 141 

Note on sodium carbonate 142 

Experiment 62. Reaction between sodium hydroxide and car- 
bon dioxide 142 

Experiment 63. Reaction between sodium carbonate and sul- 
phuric acid 143 

Test for a carbonate 144 

Experiment 64. Reaction between sodium carbonate and hy- 
drochloric acid 144 

Experiment 65. Reaction between sodium chloride and sul- 
phuric acid 145 

Definition of metathesis 146 

Experiment 66. Reaction between sodium and mercury . . 147 

Definition of amalgam . . 148 

Sodium and its compounds 149 

Experiment 67. Potassium and its properties .... 149 

Note on potassium 150 

Experiment 68. Heating potassium in contact with air . . 150 

Experiment 69. Action of water on potassium oxide . . 150 

Note on potassium hydroxide 151 

Experiment 70. The action of potassium on water . . . 152 
Experiment 71. Reaction between potassium hydroxide and 

sulphuric acid 152 

Flame coloration of potassium 153 

Flame coloration of sodium ....... 153 

Note on potassium sulphate . . . . . ' . . 154 



CONTENTS xiii 

i^The references arc to pages.) 
Experiment 72. Reaction between potassium hydroxide and 

hydrochloric acid 154 

Experiment 73. Reaction between potassium hydroxide and 

carbonic acid 154 

Note on potassium carbonate 155 

Experiment 74. Action of hot potassium on carbon dioxide . 155 

Experiment 75. Calcium and its properties .... 158 

Note on calcium 158 

Experiment 76. Heating calcium in contact with air . . 158 

Note on calcium oxide 159 

Experiment 77. Action of water on calcium oxide . . . 159 

Note on calcium hydroxide 160 

Uses of lime 161 

Experiment 78. The action of calcium on water . . . 161 
Experiment 79. Reaction between calcium hydroxide and hy- 
drochloric acid 162 

Preparation of lime-water 162 

Note on calcium chloride 164 

Experiment 80. Preparation and properties of calcium sulphate 164 

Definition of precipitate 164 

Note on calcium sulphate 165 

Further study of calcium sulphate 165 

Solubility of calcium sulphate 165 

Experiment 81. Reaction between calcium hydroxide and car- 
bonic acid 167 

Note on calcium carbonate 167 

Experiment 82. Reaction between calcium hydroxide and car- 
bon dioxide 168 

Experiment 83. Decomposition of calcium carbonate by means 

of heat 168 

Experiment 84. Reaction between calcium carbonate and hy- 
drochloric acid 171 

Best test for a carbonate 172 

Experiment 85. Reaction between calcium carbonate and sul- 
phuric acid 172 

Experiment 86. Hard water and soft water .... 173 

Definition of soft water 173 

Definition of hard water 173 

Permanently hard water 173 

Temporarily hard water 173 



xiv CONTENTS 

(The references are to pages.) 
Experiment 86 — Continued 

•Test for hard water 173 

A Chemical Investigation . . . 174 

Experiment 87. Reaction between sulphuric acid and niter, or 

the preparation of nitric acid 175 

Definition of distillation 176 

Experiment 88. The action of magnesium on nitric acid . . 178 

Experiment 89. The action of copper on nitric acid . . . 179 

Nitrogen monoxide 181 

Nitrogen dioxide .181 

Composition of nitric acid '. 181 

Experiment 90. Reaction between potassium hydroxide and 

nitric acid 182 

Theory of Chemistry 183 

Definition of Chemistry 183 

Periods in the History of Chemistry 184 

Period of Alchemy 185 

Aristotle 185 

Geber 185 

Valentine 185 

Transmutation of metals 185 

Philosopher's stone . . . 185 

Medical Period 186 

Paracelsus . 186 

Libavius 186 

Van Helmont . 186 

Glauber 186 

Period of Robert Boyle 187 

Boyle's Law . . 187 

Phlogiston Period 187 

Stahl 187 

Pneumatic Period 188 

Cavendish 188 

Black 188 

Rutherford 189 

Priestley 189 

Bergman 189 

Scheele 189 

Lavoisier 190 

Combustion theory 190 



CONTENTS xv 

(The references are to pages.) 

Pneumatic Period — Contin ued 

Law of Conservation of Mass 191 

Atomic Period 191 

Richter 191 

Bertollet 192 

Proust 192 

Law of Definite Proportions by Weight 192 

Dalton • . . 192 

Law of Expansion of Gases by Heat 192 

Tension of Aqueous Vapor 193 

Law of Multiple Proportions by Weight 193 

Atomic Theory 193 

Definition of atoms 193 

Definition of molecules 193 

Combining numbers 194 

Dalton 's symbols 194 

Gay-Lussac 195 

Law of Definite Proportions by Volume 195 

Avogadro 195 

Avogadro's Suggestion 195 

Mitscherlich 195 

Dulong and Petit 195 

Law of Dulong and Petit 195 

Prout 196 

Stas 196 

Berzelius 196 

Liebig 196 

Hofmann 196 

Woehler 196 

Bunsen 197 

Dumas 197 

Davy 197 

Faraday 197 

Faraday's Law 197 

Moissan 197 

Meyer 197 

Periodic table 197 

MendeleefT 197 

Periodic Law 197 

Language of Chemistry 198 



xvi CONTENTS 

(The references are to pages.) 
Language of Chemistry — Continued 

Symbols l 198 

Table of the elements and their symbols ... . . . 200 

Formulae 199 

Coefficients 201 

Underlying principles in writing equations 201 

Helps in writing equations 202 

Valence '. 203 

Key to the Valence of the Elements 205 

Structural formulae 206 

How to determine the formula of a salt 207 

How to determine the formula of a hydrate .... 208 
Number of atoms in the molecule of an elementary sub- 
stance 209 

Proof that the molecule of hydrogen contains two atoms . 209 

How to write an equation 210 

Writing equations illustrated 211 

List of equations for Experiments 1 to 90 . . . .212 

Experimental Work 219 

Experiment 91. An experiment to verify the Law of Conserva- 
tion of Mass 219 

Experiment 92. An experiment to verify the Law of Definite 

Proportions by Weight 221 

Experiment 93. Verification of the Law of Multiple Propor- 
tions 222 

Problems 225 

Experiment 94. Verification of the Law of Definite Propor- 
tions by Volume 226 

Experiment 95. An experiment to verify the Law of Boyle . 229 

Making a crude barometer 229 

The pressure of air 231 

Application of the Law of Boyle 233 

Problems involving the Law of Boyle . . . . . 234 

The barometer 234 

High grade barometer 235 

The vernier scale 235 

Reading the barometer 236 

Experiment 96. An experiment to verify the Law of Dalton . 236 

Absolute zero 238 

Absolute scale 239 



CONTENTS xvn 

(The references are to pages.) 
Experiment 96 — Continued 

Absolute temperature 239 

Application of the Law of Dalton 239 

Problems involving the Law of Dalton 240 

Application of both the Law of Boyle and the Law of Dalton 240 
Problems involving both the Law of Boyle and the Law of 

Dalton 241 

Experiment 97. Weight and specific gravity of air . . .241 

Definition of specific gravity 243 

Experiment 98. Weight and specific gravity of carbon dioxide 244 

Standard conditions 244 

Standard temperature and pressure 244 

Directions for using the gas balance 244 

Experiment 99. Combining number of magnesium . . . 247 

Definition of combining number 248 

Equivalent weight 248 

Correction for pressure of aqueous vapor .... 249 

Experiment 100. Determination of atomic weights . . . 250 

Definition of specific heat 252 

Application of the Law of Dulong and Petit .... 254 

Atomic weights 254 

Table of atomic weights 255 

Determination of Molecular Weights 254 

Physical method 256 

Chemical method 256 

Experiment 101. Determination of molecular weights by the 

physical method 256 

Experiment 102. An experiment to determine molecular 

weights by the chemical method 257 

Molecular weight of potassium chlorate 258 

Molecular weight of potassium chloride 258 

Atomic weight of potassium 258 

Stoichiometry 259 

Definition of stoichiometry 259 

Solution of problem in stoichiometry 259 

Rule of stoichiometry 261 

Problem 261 

Application of stoichiometry to problems involving gas vol- 
umes 261 

Problem 263 



xvni CONTENTS 

(The references are to pages.) 
Stoichiometry — Continued 

Determination of the chemical formula from the percent- 
age composition of the compound 2C3 

Problem 263 

Determination of the percentage composition of a com- 
pound from its formula 263 

Problem 264 

Experimental Work 265 

The Halogens 265 

Experiment 103. Bromine and its properties .... 266 
Experiment 104. Hydrobromic acid and bromides . . . 267 
Experiment 105. Replacement of bromine in a bromide by- 
chlorine 269 

Experiment 106. Preparation of bromine . . . . . 268 

Test for bromine 269 

Note on bromine and bromides 269 

Experiment 107. Iodine and its properties .... 270 
Experiment 108. Hydriodic acid and iodides .... 270 
Experiment 109. Replacement of iodine in an iodide by chlorine 271 
Experiment 110. Replacement of iodine in an iodide by bro- 
mine . . . ... . 271 

Note on iodine and iodides . . . . . . .271 

Experiment 111. Calcium fluoride and its properties . 272 
Experiment 112. Hydrogen fluoride; its preparation, its prop- 
erties, and its use in etching glass 272 

Note on fluorine and fluorides 273 

Experiment 113. Preparation of chlorine ... . . . 273 

Note on the Halogens 274 

Table of comparison of the Halogens 274 

Experiment 114. Arsenic and its properties .... 275 

Experiment 115. Oxidation of arsenic 275 

Experiment 116. Reduction of arsenic oxide .... 276 
Experiment 117. Arsenic sulphide; its preparation and its 

properties 276 

Test for arsenic 276 

Note on arsenic and its compounds 277 

Experiment 118. Antimony and its properties .... 277 

Experiment 119. Oxidation of antimony 277 

Experiment 120. Antimony sulphide; its preparation and its 

properties 278 



CONTENTS xix 

(The references are to pages.) 
Experiment 120 — Continued 

Test for antimony 278 

Note on antimony 278 

Experiment 121. Lead and its properties 279 

Note on lead 279 

Experiment 122. Oxidation of lead 279 

Note on the oxides of lead 280 

Experiment 123. Reduction of lead oxide 280 

Experiment 124. Action of acids on lead 280 

Experiment 125. Preparation of lead chloride by metathesis . 282 

Experiment 126. Preparation of lead sulphate bv metathesis . 283 

Experiment 127. The action of heat on lead nitrate . 283 

Definition of decrepitation 283 

Note on heating nitrates in general 283 

Experiment 128. Replacement of lead by zinc .... 283 

Experiment 129. Preparation of lead carbonate . . . 284 

Note on lead carbonate 284 

Experiment 130. Preparation of lead chromate . . . 284 

Note on lead chromate and chromium compounds . . 285 

Experiment 131. Tin and its properties 285 

Note on tin 285 

Experiment 132. Oxidation of tin 285 

Experiment 133. Action of acids on tin 286 

Experiment 134. Aluminium and its properties . . . 286 

Note on aluminium 286 

Experiment 135. Oxidation of aluminium 286 

Note on aluminium oxide 287 

Experiment 136. Action of acids on aluminium . . . 287 

Experiment 137. Alum and its properties 287 

Note on alum 288 

Note on alums in general 288 

Experiment 138. Aluminium hydrate; its preparation and its 

properties 288 

Note on insoluble hydrates ....... 288 

Experiment 139. Silver; its properties and its oxidation . 289 

Note on silver 289 

Experiment 140. Action of acids on silver 289 

Note on silver nitrate 289 

Experiment 141. Halogen salts of silver 290 

Note on photography 291 



xx CONTENTS 

(The references are to pages.) 
Experiment 142. Silver sulphide; its preparation and its prop- 
erties : 291 

Experiment 143. Replacement of silver by copper . . . 292 

Experiment 144. Reduction of silver chloride .... 293 

Nascent hydrogen 293 

Experiment 145. Silver oxide; its preparation and its prop- 
erties 293 

Experiment 146. Bismuth and its properties .... 294 

Note on bismuth 294 

Experiment 147. Oxidation of bismuth 295 

Experiment 148. The nitrates of bismuth; their preparation 

and properties 295 

Note on normal salts and basic salts 295 

Note on hydrolysis 296 

Experiment 149. Fusible alloy; its preparation and its prop- 
erties 296 

Note on alloys 297 

Composition of common alloys . 297 

Experiment 150. An experiment to prove the composition of 

ammonia 298 

Note on the composition of ammonia 301 

Experiment 151. Preparation of ammonium salts . . . 302 

Ammonium and ammonia 302 

Experiment 152. Preparation of ammonia from an ammonium 

salt 303 

Test for ammonium 303 

Experiment 153. An experiment to illustrate the great solu- 
bility of ammonia in water 303 

Experiment 154. Nitrous oxide; its preparation and its prop- 
erties 305 

Note on the oxides of nitrogen 305 

Experiment 155. Preparation of silicic acid and silicon diox- 
ide 306 

Note on silicic acid and silicon dioxide 306 

Experiment 156. Calcium silicate; its preparation and its prop- 
erties . . . 306 

Note on glass 307 

Experiment 157. Aluminium silicate; its preparation and its 

properties 307 

Note on aluminium silicate 307 






CONTENTS xxi 

(The references are to pages.) 
Experiment 157 — Continued 

Note on pottery 307 

Experiment 158. Preparation of metallic sulphides by precip- 
itation . 308 

Tests for metals 309 

Ferrous and Ferric Iron 309 

Experiment 159. Reduction of an iron solution from the fer- 
ric state to the ferrous state 309 

Experiment 160. Oxidation of an iron solution from the fer- 
rous state to the ferric state 310 

Experiment 161. The hydroxides of iron; their preparation 

and their properties 310 

Note on the oxides and hydroxides of iron . . . .311 

Note on the kinds of iron 311 

Note on the extraction of iron 311 

Note on the extraction of metals from ores in general . . 312 
Experiment 162. The reduction of copper oxide by means of 

hydrogen 312 

Note on reduction by hydrogen in general . . . .313 
Experiment 163. Calcium phosphate; its preparation and its 

properties 313 

Note on phosphoric acid and phosphates . . . .314 

Note on acid salts 314 

Experiment 164. Preparation of illuminating gas from soft 

coal 314 

Note on illuminating gas 314 

Note on water gas 315 

Note on acetylene 315 

Note on Pintsch gas 315 

Experiment 165. An experiment to determine the per cent 
of water of crystallization in crystallized barium chlo- 
ride 216 

Note on barium and its compounds 317 

Experiment 166. Tests for bases and acids in salts . . 317 

Analysis of an unknown salt 318 

Experiment 167. An experiment to illustrate Berthollet's 

Laws 319 

Experiment 168. An experiment to determine the solubility 

of a salt in water 320 

Plotting the curve of solubility for a salt .... 322 



Xxii CONTENTS 

(The references are to pages.) 

Experiment 168 — Continued 

Curves of solubility 323 

Note on saturation and kindred topics 324 

The Periodic System 325 

The periodic law „ 325 

Periodic classification of the elements . . . . . 326 

End of the Atomic Period 328 

Organic Chemistry 328 

Definition of Organic Chemistry ...... 328 

Ordinary alcohol 329 

Wood alcohol 329 

Ether 329 

Acetic acid 330 

Glucose 330 

Starch 330 

Petroleum , 330 

Soap 330 

Experiment 169. An experiment to prove the presence of car- 
bon and of hydrogen in certain organic compounds . 331 

The Modern Period 332 

Faraday's Law 332 

Arrhenius 333 

Solution and dissociation 333 

Application of electrolysis 335 

Ostwald 335 

Dewar 335 

M. and Mme. Curie 335 

Ramsay 335 

Raoult 335 

Thermochemistry 336 

Pfeffer 336 

Van't Hoff 336 

Reversible Reactions 337 

Chemical Equilibrium . . . 337 

Law of mass action . , . . 338 

Gibbs 338 

The phase rule 338 

Conclusion 338 

Outline of First Year's Work 339 

Appendix 355 



CONTENTS xxiii 

(Thereferences are to pages.) 

Appendix — Continued 

Table of elements with some of their properties . . . 356 

Additional problems for class use 362 

Practical questions 374 

List of text-books for reference library 377 

List of individual apparatus 378 

List of general apparatus 379 

List of chemicals 281 

Suggestions to the teacher 382 

Answers to additional problems 388 

Index 391 



ILLUSTRATIONS 

Fig. Title Pag 

1 Various articles mentioned in the list of apparatus opposite 1 

2 Reading the graduate 7 

3 Holding the graduate correctly for reading .... 7 

4 Platform balance . 8 

5 Set of iron weights 9 

6 Horn-pan balance 10 

7 Set of smaller weights 11 

8 Bunsen burner 13 

9 Bunsen burner in detail 13 

10 How to break glass tubing 14 

11 Heating glass tubing for bending 15 

12 Bending glass tubing 15 

13 A good bend 16 

14 A poor bend 16 

15 The wrong way to draw out a glass tube .... 16 

16 A poor tip 16 

17 The right way to draw out a glass tube .... 17 

18 A good tip 17 

19 The proper way to hold the tube while it is cooling . . 17 

20 The best shape for a rubber connector 18 

21 A finished wash-bottle 19 

22 Theoretically perfect crystals 20 

23 Single crystals as found in nature 20 



xxiv CONTENTS 

Fig. Title Page 

24 Single crystals as found in nature and as made in the lab- 

oratory 20 

25 Large clusters of crystals as found in nature ... 20 

26 Blast lamp 23 

27 Interior construction of the blast lamp . . . . .23 

28 Foot bellows 24 

29 Richards blower 25 

30 Heating copper turnings 26 

31 Using the test tube holder .31 

32 Determining the proportion of oxygen in air ... 34 

33 Glass tube with end fallen together just a little ... 44 

34 Glass tube just ready to close 44 

35 Finished bulb tube 44 

36 Testing a gas with a glowing splinter 44 

37 Apparatus for generating oxygen on a large scale . . 47 

38 Transferring a dry powder to a flask 48 

39 First step in folding a filter 50 

40 Second step in folding a filter 50 

41 Third step in folding a filter 50 

42 Fourth step in folding a filter ... ... 50 

43 A finished filter ...".. 50 

44 Filtering 57 

45 A very simple apparatus for electrolyzing water ... 57 

46 A convenient electrolysis apparatus 58 

47 Hand dynamo 59 

48 Plunge batteries . .59 

49 Apparatus for showing the volumetric composition of water 61 

50 Apparatus for passing steam over magnesium ... 66 

51 Testing the gaseous product for inflammability ... 68 

52 Snow crystals 70 

53 Snow crystals .70 

54 Catching oxygen by displacement of air .... 74 

55 Catch bottles 79 

56 Apparatus for making the second oxide of sulphur . . 80 

57 The generator ready for general use ..... 86 

58 The generator set up for generating hydrogen ... 87 

59 Apparatus for passing hydrogen into warm sulphur vapors 109 

60 Apparatus for passing carbon oxide over hot zinc . . 120 

61 Test tube generator 127 

62 Apparatus for treating sodium with water . . . .135 



CONTEXTS xxv 

Fig. Title Pag e 

63 The proper way to use litmus paper 138 

64 Apparatus for catching a small quantity of gas . . . 148 

65 Apparatus for passing carbon dioxide over potassium . 156 

66 Apparatus for decomposing calcium carbonate by means of 

heat 169 

67 Apparatus for decomposing niter with sulphuric acid . . 175 

68 Bunsen 196 

69 Some more noted chemists 196 

70 Some more noted chemists 196 

71 Mendeleeff 196 

72 Steam bath 221 

73 A crude barometer 230 

74 The Boyle tube 232 

75 A high grade barometer 235 

76 The vernier 236 

77 The absolute scale 239 

7S The gas balance 245 

79 A photographic positive (print) 291 

80 A photographic negative (plate) . . . . . .291 

81 Apparatus for proving the composition of ammonia . 300 

82 The ammonia fountain 304 

53 Apparatus for reducing copper oxide with hydrogen . .313 

54 A sand bath in use 320 

85 A pipette in use 321 

86 Plotting the curve of solubility of a salt .... 322 

87 Curves of solubility of some common salts .... 323 



FIRST YEAR CHEMISTRY 



Introduction. — This text is so written that the student 
can begin his laboratory work with the following paragraph. 

Apparatus. — Each student should be supplied with the 
following pieces of apparatus, all of which are necessary for 
the work outlined in this book. The articles marked with 
a star are illustrated in the sketch on the opposite page. 



Locker Key No 

Text-book 

Laboratory notebook 

Ring stand and three rings* 

Two clamps for stand* 

Tripod* 

Four fruit jars and washers* 

Three prescription bottles* 

Two catch bottles* 

Nest of three beakers* 

Three flasks* 

Kjeldahl flask* 

Graduated cylinder* 

Two funnels* 

Thistle tube* 

Twelve test tubes* 

One stick glass rod 

Two sticks glass tubing 

Mortar and pestle* 

Porcelain evaporating dish* 

Porcelain crucible and cover* 



Nest of four Hessian crucibles* 

Test tube rack* 

Test tube brush* 

Bunsen burner and hose* 

Bat- wing burner and hose* 

Box of weights* 

Twenty round filters 

Ten rubber connectors 

Ten assorted corks 

Round file 

Triangular file 

Iron rod 

Brass forceps* 

Crucible tongs* 

Two iron gauzes 

Tubing clamp* 

Deflagrating spoon* 

Pipe-stem triangle* 

Horn spatula* 

Thirty centimeter rule 

Two towels 



Laboratory Notebook. — Each student should be supplied 
with a notebook, in which to record the results of his labora- 
tory work. The most convenient book to use is one hav- 
ing about four hundred pages of unruled paper, with pages 
seventeen by twenty-one centimeters in size. The book 
should be strengthened with leather back and corners, be- 



2 FIRST YEAR CHEMISTRY 

cause laboratory usage is apt to destroy the light cloth back 
of the ordinary pasteboard-covered blankbook. 

No descriptive text-book is used in the course. Since 
the student gets his knowledge of chemical substances thru ac- 
tual contact with the substances themselves, the laboratory note- 
book becomes a storehouse for his observations from the experi- 
ments. Therefore, he should keep his records of experi- 
ments in such a way that he can use the notebook as his 
text-book later in the year. Two full pages that face each 
other should be reserved for each experiment. On the left- 
hand page should be put the first rough pencil notes at the 
moment the observations are made; here also should ap- 
pear all the figuring called for by the experiment. Note- 
taking and figuring on scraps of paper, or anywhere but in 
the notebook, always lead to waste of time and often to ihe 
necessity of repeating the experiment, so easy is it for loose 
papers to get lost. In case of mistakes in numerical work 
do not correct by erasing and rewriting, but cross out the 
old and let the new appear by its side. The left-hand page 
is reserved for such preliminary notes, and such corrections 
do not spoil the notebook in the estimation of those who 
know for what purpose the notebook is designed. It has 
been found best to use the left-hand page only for notes 
taken at the time of the experiment and for all calculations. 
Such pencil notes need not contain a full account of the ex- 
periment, but simply those statements needed to answer 
questions asked in the text. If, however, one page happens 
to be not enough for the notes, continue the pencil notes on 
the next left-hand page of the book. On the right-hand 
page should be written a good and full account of the ex- 
periment; this account should be in ink and should be 
written as soon as possible after the work has been re- 
viewed in class; furthermore, it should reveal to the reader 
what was done and observed; also what conclusion, if any, 



FIRST YEAR CHEMISTRY 3 

the student reached in regard to the experiment. Be sure 
to answer all questions asked in the text, and to mark all nu- 
merical results so that they can be found easily when wanted. 

All new pieces of apparatus, the crystalline form of new 
compounds, and anything else about the experiment that 
can be so represented, should be represented by sketches. 
A drawing often saves one the necessity of making long 
descriptions of apparatus; a few lines, even tho crudely 
drawn, often serve the purpose as well as elaborate descrip- 
tions. When making a sketch always have the apparatus 
before you. and show as far as possible the relative sizes of 
the different parts of the apparatus. It is much better to 
make drawings directly from the apparatus than to copy 
the illustration in the book; it is allowable, however, to 
refer to the figures in the text to learn how to convey the 
visual picture of the apparatus to paper. Every student, 
even if he is not artistic, can make some kind of a drawing; 
his sketch should be at least so near like the original that 
a person not familiar with the article could pick one out 
of the stock room after inspecting the sketch. Drawings 
made in pencil are generally satisfactory, unless one is apt 
with ink. The drawing should be made but once, and on 
the right-hand page. 

The notebook should always be kept in the laboratory; 
it must not be taken from the laboratory even to write up 
the ink records, except by permission from the instructor; 
after a little practice the ink record of experiments already 
reviewed in class can easily be made while waiting for some 
new experiment to run to completion. The book should 
always be brought to class, and once a week it should be 
handed in at the instructor's office to be examined. 

The Preliminary Work, which is covered by the first four 
experiments, is designed to lead up to the experiments of 
a strictly chemical nature. The first two experiments fa- 



4 FIRST YEAR CHEMISTRY 

miliarize the student with the metric units used during the 
year; the third experiment explains the use of the Bunsen 
burner and gives practice in bending and otherwise manip- 
ulating glass tubing; the fourth tells how to make a wash- 
bottle. 

Experiment 1. Measuring. Have ready the following 
apparatus: A rule, about a foot long, with inches divided 
into eighths on one side, and centimeters divided into milli- 
meters on the other side; a glass cylinder holding 100 cubic 
centimeters and graduated to cubic centimeters; three pre- 
scription bottles, one large, one medium, one small; three 
lip beakers, large, medium, and small; three flasks, large, 
medium, and small; and a test tube rack containing the 
three different sizes of test tubes. 

The Metric System. — As the measurements made in sci- 
entific work are always taken in the metric system, it is 
well to have enough practice with the metric units to get a 
working acquaintance with them. Complete metric tables 
may be found in any arithmetic, but for our work the fol- 
lowing are sufficient; 

LINE OR LINEAR MEASURE. 

10 millimeters (mm.) = 1 centimeter (cm.) 
10 centimeters = 1 decimeter (dm.) 

10 decimeters = 1 meter (m.) 

SURFACE OR SQUARE MEASURE. 

100 square millimeters (sq. mm.) = 1 square centimeter (sq. cm.) 
100 square centimeters = 1 square decimeter (sq. dm.) 

100 square decimeters = 1 square meter (sq. m.) 

VOLUME OR CUBIC MEASURE. 

1000 cubic millimeters (cu. mm.) = 1 cubic centimeter (cu. cm.) 
1000 cubic centimeters = 1 cubic decimeter (cu. dm.) 

1000 cubic decimeters = 1 cubic meter (cu. m.) 



FIRST YEAR CHEMISTRY 5 

MASS OR WEIGHT MEASURE. 

10 milligrams (mg.) = 1 centigram (eg.) 

10 centigrams = 1 decigram (dg.) 

10 decigrams = 1 gram (g.) 

1000 grams = 1 kilogram (kg.) or kilo (K) 

Square measure finds but little use in chemical work. The 
linear units, — generally the centimeter, sometimes the milli- 
meter, — are used for measuring glass tubes and other parts 
of apparatus. Cubic measure is used extensively, glass 
vessels being designated according to their capacity in cubic 
centimeters; liquids are usually measured in cubic centi- 
meters; on account of the extensive use of this unit of 
volume the shorter abbreviation, c.c, is generally used for 
the cubic centimeter instead of the longer abbreviation, 
cu. cm., given in the table above. Larger volumes of 
liquids are often measured by the cubic decimeter, which 
is then called a liter and is abbreviated to L. In the labo- 
ratory, weights of substances are generally expressed in 
grams; for accurate work the fractional parts of the gram 
are used, and in hea\T work the kilo is sometimes used. 

Measurement of desk. — Measure, in the metric system, 
the length of your laboratory desk; also its width. It will 
be found most convenient if all measurements are taken, 
both here and elsewhere during the year, in centimeters 
and afterwards reduced to any other denomination desired. 
In recording measurements use the decimal point, e.g., if 
the length of your desk is five decimeters, seven centimeters, 
and six millimeters, do not record the measurement 5 dm., 
7 cm., and 6 mm., but record it 57.6 cm. or 0.576 m., according 
i intend to express the length in centimeters or in meters. 

Be sure to make the pencil record of these measurements 
on the left-hand page of the notebook. Also let all calcu- 
lations appear on this same page. It is well to label answers 
so that they ran be found easily when wanted. 



6 FIRST YEAR CHEMISTRY 

From the length and width, both expressed in centi- 
meters, calculate the area of the desk top, first in square 
centimeters, then in square millimeters, next in square 
decimeters, and finally in square meters. 

Now measure the length and width of the desk again, 
this time in the English system; read to the nearest eighth 
of an inch, and record the measurements in inches and 
eighths of an inch. Calculate the area of the desk top in 
square inches. 

Since the cumbersome English measures have not been 
entirely replaced yet by the more convenient metric meas- 
ures it is well to know the relation between the units of the 
two systems. From the length of the desk, expressed in 
centimeters and in inches, find out how many centimeters 
there are in one inch; then from this answer calculate how 
many inches there are in one meter. From the area of the 
desk expressed in surface units of both systems find out how 
many square centimeters there are in one square inch, and 
roughly what part of a square inch a square centimeter oc- 
cupies. Fix in mind the round numbers for the answers 
obtained in this paragraph. 

What shorter method is there that, from the materials 
at hand, will give the number of centimeters in an inch? 
Why does this easier method not give so accurate an an- 
swer as the method actually used? 

Measurement of glass vessels. — Take the large glass cylin- 
der, graduated in cubic centimeters and commonly called 
the " 100 c.c. graduate"; fill the largest beaker with water, 
not brim full but conveniently full only; pour the water 
from the beaker into the graduate and find out how many 
cubic centimeters of liquid you could comfortably work 
with in the large beaker. Find the capacity of the medium 
and small beakers in the same way you found the capacity 
of the large beaker. Find also the capacity in cubic centi- 



FIRST YEAR CHEMISTRY 




Fig. 2. 
Reading the graduate. 



meters of the three prescription bottles, of the three flasks, 

and of three test tubes of different sizes. 
How to read the graduate. — -Reading the graduate should 
be done with some care. When the 
surface of the water in the graduate 
is examined from the side it may 
look as if there were two surfaces 
instead of one. This is due to the 
fact that the water is drawn up on 
the sides of the glass a little as 
shown in Fig. 2. The main surface 
of the water is at AA and this, of 
course, is the level that should be 
read. The small amount of water 
drawn up against the sides of the 
glass above the level A A may be 
neglected. The correct reading of 

the graduate in Fig. 2 is, therefore, 

4-4 c.c. Naturally the graduate 

must be held exactly vertical and 

the eye should be on a level with 

the surface A A . This is most easily 

accomplished by holding the cyl- 
inder near the top and raising 

it till the surface of the water is 

just opposite the eye, as shown 

in Fig. 3. If the eye is much 

below or much above the line A A, 

the reading will be inaccurate, as 

shown by the lines BB and CC in 

Fig. 2. 

If the exact volume of water in the beaker is desired, 

care should be taken to shake out all water from the grad- 
uate before filling it again. Careless emptying of the gradu- 




Fig. 3. Holding the graduate 
correctly for reading. 



s 



FIRST YEAR CHEMISTRY 



ate often leaves a whole c.c. or more of water clinging to 
the inside of the graduate, and if the volume to be meas- 
ured is large that error will then be multiplied as many 
times as the graduate is filled. In careful quantitative 
work such an error would render the experiment valueless. 

One purpose of this work in measuring is to enable you 
to tell at a glance about how much the laboratory vessels 
in common use contain, so that if a new experiment calls 
for 130 c.c. of liquid you can tell without hesitation whether 
a test tube or a beaker will be most suitable. As far as pos- 
sible fix the rough capacities of the various vessels in mind. 



Experiment 2. Weighing. This experiment is similar 
to the one on measuring, except that the gram instead of 
the cubic centimeter is used as the unit of measure. 

Have ready a platform balance, a set of iron weights, a 
horn-pan balance, a set of small brass weights, brass for- 
ceps, 100 c.c. graduate, small beaker, three large nails, and 

any three of the larger 
glass vessels. 

Before proceeding with 
the actual weighing indi- 
cated below, examine the 
balances and weights care- 
fully to make sure that 
they are in order and that 
you know how to use them. 
The following paragraphs 
will be found helpful in 
this examination. 

The platform balance 

Fig. 4. Platform balance. should be Ca P able ° f hold - 

ing a maximum load of 5 

kilograms and should be sensitive to a tenth of a gram, 




FIRST YEAR CHEMISTRY 



9 



i.e., it should show a variation of one tenth of a gram 
in the weight of an article being weighed. This bal- 
ance is pictured in Fig. 4. First see that the rider is on 
the zero mark of the beam; then see if the balance is in 
equilibrium, i.e., if it swings easily when set in motion, and 
if the pans are at the same level when they are at rest. In 
case the balance does not swing easily, see whether the 
knife edge bearings are in their proper places. If the bear- 
ings are in place and the balance still sticks, perhaps a drop 
of oil on the bearings will remedy the trouble. If the pans, 
when swinging properly, do not come to rest at the same 
level, move the equilibrium nuts on the threaded rod in 
the middle of the balance a little to the right or left, or 
put some small pieces of metal, wood or paper on the lighter 
pan till equilibrium is gained. 

The set of iron weights, 10 grams to 2000 grams inclu- 
sive, should always be used with the plat- 
form balance. This set, shown in Fig. 5, 
should contain the following pieces: 



10 gram 
20 gram 
20 gram 
50 gram 
100 gram 



100 gram 

200 gram 

500 gram 

1000 gram 

2000 gram 




Fig. 5. 

Set of iron weights. 



If the set is complete all the pieces smaller 
than one kilo should together amount to 
1000 grams. 

How to weigh. — Always put the article to 
be weighed on the left-hand pan and the weights on the right- 
hand pan. Begin with a weight that is plainly too heavy. 
Suppose it is the 1 kilo. If your guess was right, remove the 
weight that was too heavy and try the 500 gram weight. If 
this is too light add the 200 gram weight; if still too light add 



10 



FIRST YEAR CHEMISTRY 



a 100 gram weight; if still too light add the other 100 gram 
weight. If still too light it shows that the article in ques- 
tion weighs between 900 and 1000 grams. Next add the 
50 gram weight. If this is too heavy remove the 50 gram 
weight and put on a 20 gram weight; if still too light put 
on the other 20 gram weight. If it is now too light it shows 
that the article weighs between 940 and 950 grams. To 
get still closer to the true weight move the rider from the 
zero mark along the beam towards the right till equilibrium 
is established; moving the rider from one end of the beam 
to the other is equivalent to adding 5 grams to the weight 
pan; the beam is graduated so that by moving the rider 
from mark to mark it is possible to read to grams and 
tenths of a gram. If equilibrium occurs with the rider on 
4.5 the article weighs 944.5 grams. 

The horn-pan balance should 
be sensitive at least to one one- 
hundredth of a gram. This bal- 
ance is for delicate weighing and 
should never be loaded with 
more than 150 grams; it is pic- 
tured in Fig. 6. See that the 
balance moves easily on its bear- 
ings, and that it is in equi ib- 
rium; if necessary, add some 
pieces of fine iron wire or paper 
to either pan to bring the pointer 
to the center. With this bal- 
ance, as with the platform bal- 
ance, articles to be weighed 
should be put in the left-hand 
pan and the weights in the right- 
hand pan. 
The set of smaller weights, 10 milligrams to 100 grams 






Fig. 6. Horn-pan balance. 



FIRST YEAR CHEMISTRY 11 

inclusive, should always be used with the horn-pan balance. 

They are kept in a wooden box as shown in Fig. 7. These 
weights, from 1 gram to 100 grams, 
should be of brass and of these values: 




1 gram 


5 gram 


20 gram 


2 gram 


10 gram 


50 gram 


2 gram 


10 gram 


100 gram 



Fig. 7. 

weights. All the brass weights below 100 grams 

should together amount to 100 grams. 
The fractional parts of a gram should be of these values: — 

10 milligram 100 milligram 

20 milligram 100 milligram 

20 milligram 200 milligram 

50 milligram 500 milligram 

If the set is complete these fractional parts of a gram should 
together equal just one gram. All the weights of this small- 
er set should be handled with brass forceps. 

Heavy weighing. — Fill three of the larger vessels con- 
veniently full of water; weigh each separately on the plat- 
form balance as described above. Weigh accurately to a 
single gram, i.e., have the weight you record within one 
gram of the true weight. Record in }^our notebook the 
weights of the three vessels in grams and decimal parts of 
a gram, using the decimal point as in the case of the linear 
measurements. Then weigh all three vessels together and 
see if the weight of the three together equals the arithmetic- 
al sum of the three when weighed separately. 

Light weighing. — Using the horn-pan balance and the 
set of smaller weights, weigh three large nails separately. 
Weigh accurately to centigrams, i.e., have the weight you 
record within one centigram of the true weight. Record 
the weight in grams and decimal parts of a gram. Then 



12 FIRST YEAR CHEMISTRY 

weigh all three nails together and see if the weight of the 
three together equals the arithmetical sum of the three 
when weighed separately. The difference between the two 
should be not more than 0.10 gram. If the variation is great- 
er, repeat all the weighings. 

Tare weighing. — Put the large graduate empty on the 
left-hand pan of the platform balance; on the right-hand 
pan put a small beaker and add water to it till equilibrium 
is reached. The beaker and water are called a " tare weight," 
or simply a "tare," and the actual weight of it is not taken 
into consideration; the graduate is said to be " balanced 
with a tare." Then, leaving the empty graduate on the 
left-hand pan and the tare on the right-hand pan, fill the 
graduate with water to the 100 c.c. mark. Add the neces- 
sary weights to the right-hand pan to get equilibrium 
again; these weights give directly the weight of the 100 c.c. 
of water. The weight of the water should, of course, be 
just 100 grams, since the metric system was so devised that 
one gram should be the weight of one c.c. of water. If 
your result is not consistent with this fact, try to find out 
what caused the discrepancy. 

Experiment 3. Manipulating glass. The ability to ma- 
nipulate glass tubing is essential to success in chemical 
work. The following paragraphs explain how to cut, fire- 
polish, bend, and draw glass tubing, these operations being 
the ones most commonly used. 

Have ready a Bunsen burner, bat-wing burner, triangular 
file, and several sticks of soft glass tubing. Then proceed 
with the directions in the following paragraphs. 

The bat-wing burner, sometimes called the fish-tail burn- 
er, is the ordinary illuminating gas burner so long used be- 
fore the introduction of the Welsbach. The gas escapes 
thru a narrow slit in the clay tip; see Fig. 11; it burns with 



FIRST YEAR CHEMISTRY 



13 






C 



£> 



a yellow flame because all the carbon in the gas has not 
an opportunity to burn completely. The unburned carbon, 
heated to luminosity or incandescence, causes the lumin- 
ous or yellow flame; it deposits itself as soot on any cold 
object held in the flame. In the laboratory this burner is 
used for bending glass tubing. 

The Bunsen burner, named after the 
German chemist, Robert Bunsen, who de- 
vised it, is a modification of the bat-wing 
burner and allows complete combustion 
of the gas thru the admixture of an ex- 
cess of air. The burner is shown com- 
plete in Fig. 8; the q 
separate parts are 
shown in Fig. 9. The 
— ift gas enters the base 
n B at the side tube A 
and escapes thru the 
small opening in the 
top of the plug C into 
the longtubeZ). Near the lower end of the 
tube D are two holes E opposite each 
other; the collar F that fits around the 
long tube has holes opposite those in the 
tube D. When the gas rushes out of the 
small opening in the top of C it draws 
air in thru the holes E and the diluted 
gas burns at the top of the burner G 
with a bluish or nearly colorless flame. 
This flame, which should be 10 to 15 
centimeters high, consists of two cones, 
an inner colorless cone of unburned gases 
and an outer or bluish cone where the 



Fig. 8. 
Bunsen burner 



O 



C 



^ 



Fig. 9. 

Bunsen burner in 
detail. 



gas is entirely consumed. The hottest part of the flame is 



14 



FIRST YEAR CHEMISTRY 



near the top and the temperature here is in the neighborhood 
of 1000° Centigrade. Light the burner by turning on the gas 
full and holding a lighted match about 5 centimeters above 
the top of the burner. Regulate the size of the flame by 
turning the gas stopcock. If the flame is not colorless 
turn the collar F till the yellow disappears from the flame. 
Always use the colorless Bunsen flame, unless specially 
directed otherwise. Occasionally when the flame is turned 
low, or when a strong draft plays upon the burner, the 
flame will "snap back" and burn at the tip C. The flame 
then has a disagreeable odor, deposits soot, often burns 
green, and sometimes emits a slight whistling noise. To 
remedy this snapping back, turn off the gas and light the 
burner anew. In the laboratory this burner is used for 
firepolishing and for drawing out glass tubing, also for gen- 
eral heating purposes. 

How to cut glass tubing. — Glass rods and tubes of small 
diameter may be cut as follows: With a sharp triangular 
file make a scratch at the point where the tube is to be 

cut. One decisive move- 
ment of the file across 
the glass in a direction 
away from you is much 
better than sawing back 
and forth lightly with the 
file, for the scratch is 
then sharp and clean, in- 
stead of indefinite and 




Fig. 10. 



How to break glass tubing 

ragged. Take the tube in the hands with the two thumb 
nails close together and touching the glass on the side op- 
posite the scratch. Spread the fingers of each hand out on 
the tube to the right and the left of the scratch as shown 
in Fig. 10. A slight pressure with the little fingers should 
cause the tube to snap off squarely at the mark. 



FIRST YEAR CHEMISTRY 



15 




How to firepolish glass tubing. — Freshly cut glass tubes 
are always sharp on the ends. These sharp edges should 
always be smoothed off by holding the end of the tube 
in the top part of the colorless Bunsen flame, constant- 
ly turning the tube between the fingers. The glass soon 
becomes soft and rounded. Care should be taken not to 
heat the tube too much, for then the glass might fuse so 
much as to constrict the diameter of the tube, or even to 
close the end of it. 

How to bend glass tubing. — The bat-wing burner, and 
not the Bunsen burner, must be used for this purpose. For 
bending tubes 
that are 6 mm. 
or less in outside 
d iameter, the 
flame should be 
a b o u t 6 cm. 
across at the wid- 
est part; for 
tubes of a larger diameter use a larger flame. Grasp 
the tube lightly between the thumb and fingers of each 

hand, the knuck- 
les of ,the left 
hand being up 
and those of the 
right hand down. 
Hold the tube, at 
the place where 
it is to be bent, 
in the upper, lu- 
minous, and hot- 
test part of the 
Fig. 12. Bending class tubing. , , a 

broad flame as in 

Pig. 11; hold it Lengthwise in the flame so that at least 



Fig. 11. Heating glass tubing for bending. 




16 FIRST YEAR CHEMISTRY 

6 cm of the tube become hot. It will probably become 
covered with soot, but that can be wiped off when the tube 
has cooled. As soon as you have put the tube into the 
flame, rotate it slowly and constantly by turning it with 
the thumb and fingers. The tube is more likely to be 
heated evenly if it is rotated continuously in one direc- 
tion than if it is simply rolled forward and backward be- 
tween the fingers. When the glass tube has softened so 
that it feels flexible, remove it from the flame, and 

1 then bend it, deliberately r — 

but not too slowly, to the ex- 
act angle wanted. See Fig. 
12. Lay it aside to cool on 
some object that will not 
crack the hot glass. If the 
Fig. 13. first attempt does not give the FlG - 14 - 

A good bend. desired ^d, take a new piece A poor bend ' 
of glass and start again. The bend should have an even curve. 
Fig. 13 shows such a bend. Fig. 14 shows the bend one 
gets by using the Bunsen burner or by using the bat-wing 
flame crosswise. 

How to draw out glass tubing. — Hold the tube between 
the thumbs and fingers, as directed in the preceding para- 
graph, but in the 

j z>=<r " ] to P P art °f the 

Bunsen flame. 
Fig. 15. Wrong way to draw out a glass tube. ^urn the tube 

continually while it is in the flame. The tube should be 
softened more for drawing than for bending. When the 

tube has softened so much that it • ^ 

begins to sag, remove it from the 
flame and gently pull the two ends IG ' ' poor ip ' 
apart. The result is shown in Fig. 15. By cutting at A a 
stubby tip like that in Fig. 16 is produced. A much better 



FIRST YEAR CHEMISTRY 



17 



drawn out tube is obtained as follows: Heat the tube as 
before till very soft, remove it from the flame, draw the 

ends apart just 

I ^- C — 1 a little, let the 

Fig. 17. The right way to draw out a glass tube. soft g lass " set " 

a couple of sec- 
onds, and then rapidly draw out more and get Fig. 17. 

This, when cut, produces the _ > 

Fig. 



of 



18. 



c a pillar v tip 

The capillary part may be FlG " 18 " A ^ ood ti P- 

kept straight while the glass is cooling and hardening by hold- 
ing the tube in a vertical position as shown in F g. 19. 



Experiment 4. Making 




Fig. 19. 

ifr way to hold the 



The proper way to hold 
tube while it is cooling 



a wash-bottle. Have ready a 500 c.c. 
flask, a good, sound cork to fit it, 
an iron rod, triangular file, round 
file, Bun sen burner, bat-wing 
burner, rubber connector, piece of 
filter paper, and several sticks of 
glass tubing. 

To the 500 c.c. flask fit a good 
cork that is free from cracks and 
air passages. Put the cork on 
the floor and with the sole of the 
shoe roll it to soften it and make 
it fit the neck of the flask tightly. 
About one third of the cork 
should protrude above the top 
of the flask. With an iron rod 
that has been held in the Bunsen 
flame a moment burn two holes 
lengthwise thru the cork. Trim 
these holes out with the round file 
till a piece of wash-bottle tubing 



18 FIRST YEAR CHEMISTRY 

can just be pushed in. The holes should be true and round, 
so that no air passages are left beside the glass. They 
should also be perpendicular to the flat ends of the cork 
and far enough apart so that the cork will not be broken out 
between the holes when the glass tubes are inserted later. 
Cut off a piece of glass tubing 10 cm. longer than the 
hight of the flask. At a point 5 cm. from one end bend 
this glass tube so that the shorter limb forms an angle of 
45° with the other. An aid in bending a glass tube to a 
definite angle is to hold the tube, just after it has been 
bent but before it has had time to harden, over two pencil 
lines drawn to the desired angle; for most purposes, how- 
ever, it is sufficient to bend the glass "by the eye" only. 
Pass the longer limb down thru one hole of the cork, and 
at a point half way from the cork to the end of the longer 
limb bend the tube a little, just enough so that the lower 
end of the tube, when the cork is inserted in the flask, shall 
be at the vertex of the angle made by the bottom of the 
flask with its side. Be sure to make the second bend to- 
ward the first. Cut off a piece of glass tubing 15 cm. long, 
and bend this at the middle to form an obtuse angle of 
about 135°. Cut off another piece of glass tubing 12 cm. 
long and draw it out at the middle to form a capillary tip. 
Snap it off at the drawn out part so that 2 or 3 cm. of ca- 
pillary are left on the part to be used for the tip. Wipe the 
soot from all the tubes and firepolish all the ends. Insert 
the second bent tube in the other hole of the cork, in line 
with the first tube 
but pointing away \ ^ ^^"1 \ " ~x^ 1 

from it, and flush with __ „ " 

, . , . . ,,, Fig. 20. Best shape for a rubber connector. 

the under side of the 

cork. With a rubber connector, — a piece of rubber tube 

about 5 cm. long, — attach the tip to the short limb of the 

first tube. The connector should fit tightly. If the rubber 



FIRST YEAR CHEMISTRY 



19 



connector is cut diagonally at the ends, as shown in Fig. 20, 
it is easier to slip the glass tubes into it. 

Fill the flask to the neck with distilled water and insert 
the cork with its fittings. The finished wash-bottle should 
look like Fig. 21. The mouth 
piece and the tip should be in as 
nearly a straight line as possible. 
The tip forms a small stream of 
water, which, owing to the flexible 
connection of rubber, can be di- 
rected by the fingers in different 
directions when the wash-bottle 
is used in later experiments. 
The slight inner bend in the longer 
tube allows all the water to be 
blown from the bottle when the 
flask is tipped up, as it generally 
is when used. Distilled water 
must always be used in the wash- 
bottle, because tap water often 
contains considerable iron rust 
and other sediment that coats 
the flask and clogs the tip. Keep the wash-bottle always 
set up and filled with distilled water ready for use. 




Fig. 21. Wash-bottle. 



Experiment 5. Copper and its properties. Take copper 
in its various forms, — sheet, turnings, foil, and wire of two 
sizes, coarse and fine, — and examine them for the proper- 
tic- mentioned in the following paragraphs. 

Definition of property. — A property of a substance, in its 
chemical meaning, is any quality which is peculiar to that sub- 
stance and which enables one to distinguish it from other 
mbstances. For instance, the whiteness of snow, the heavi- 
ness of lead, the sweetness of sugar, the transparency of 



20 FIRST YEAR CHEMISTRY 

glass are all called properties of the substances mentioned. 
Since a substance is identified by means of its properties, 
it is to our advantage to learn as many properties as possi- 
ble of each substance we meet or make during the year; 
and these properties are best learned by personal observa- 
tion and examination. 

List of properties. — The properties generally looked for 
are given in the following list: 

State Malleability Solubility 

Crystalline form Ductility Specific gravity 

Color Transparency Fusibility 

Luster Opacity Inflammability 

Hardness Translucency Boiling 

Tenacity Taste Volatility 

Brittleness Odor Consistency 

In many substances some of these properties will naturally 
be wanting, but they should all be looked for. Occasion- 
ally a substance may possess some marked property not 
mentioned in the above list; all such special properties 
must, of course, be recorded. 

How to test for properties. — The statements in this para- 
graph are given with special reference to copper, but they 
are applicable to all substances; therefore, in case of doubt 
at any time, refer to this paragraph. By state is meant 
the existence of the substance as a solid, a liquid, or a gas; 
if it is a solid, it is well to describe the form in detail, i.e., 
tell whether the substance is in lumps, sheet, wire, powder, 
or mossy; a solid usually has a definite form or shape, a 
liquid has no definite form, but takes the shape of the ves- 
sel that contains it so far as it fills that vessel, while a gas 
has no definite form but tends to spread out and fill com- 
pletely the containing vessel. A substance that has crys- 
talline form generally exists in separate fragments, each 




Fig. 22. Theoretically perfect crystals. A, octahedron; B, cube; 
C, prism; D, terminated prism. 




Fig. 23. Single crystals found in nature. A, gypsum; B, beryl; 
C, feldspar; D, amethyst. 




Fig. 24. Single crystals as found in nature and as made in the labo- 
ratory. A. magnetite; B, blue stone; C, niter; D, alum (octahe- 
dron;; E, alum (modified octahedron); F, corundum; G, pyrites. 




Fig. 25. Large clusters of crystals as found in nature. A, quartz: 
B, fluorspar. 



FIRST YEAR CHEMISTRY 21 

one having a definite geometrical form with straight edges 
and flat, lustrous faces. Many of the crystalline substances 
studied during the } r ear are prismatic, cubical or octahe- 
dral; Figs. 22 to 25 show crystals. The study of crystals 
will be taken up more in detail later when we come to 
those substances that have good crystalline form. The 
color on the outside of a substance is often not the true 
color; the true color may best be ascertained by scraping 
off the outside coating with a knife or file. The luster, or 
shine, sometimes appears to best advantage on the freshly 
cut surface; the permanency of the luster also should be 
noted; sometimes the luster remains unchanged tho the 
substance be exposed to air for a long time, sometimes it 
changes slowly during exposure for a few days, sometimes 
it changes so rapidly that the fresh surface retains the lus- 
ter only a few minutes. For hardness the substance should 
be compared with other substances; generally the harder 
of two substances wil scratch the softer; cutting with a 
knife is also a good test. Brittleness is indicated by break- 
ing into irregular fragments; it varies from crushing between 
the fingers to breaking under the hammer on the anvil. 
Tenacity means a tendency to hold together, and it may 
be tested by trying to pull the substance apart. Mallea- 
bility may be inferred from the existence in sheet form; 
it may also be shown by pounding the substance to a thin 
sheet on the anvil. Ductility may be inferred from the ex- 
istence of the substance in wire form; substances not duc- 
tile cannot be drawn nto wire. Transparency: a trans- 
parent substance is one thru which light passes and thru 
which objects may be distinguished with ease. Opacity : 
an opaque substance is one thru which light does not pass 
and thru which objects cannot be seen. Translucency : a 
translucent substance is half way between a transparent 
body and an opaque one, i.e., light passes thru it but ob- 



22 FIRST YEAR CHEMISTRY 



jects cannot be seen thru it. By taste is meant, of course 
the impression made upon the nerves of the tongue by con 
tact between the tongue and the object in question; the 
following tastes may be identified easily: sweet, sour, salty, 
bitter, puckery, and metallic; to try the taste of anything 
it is sufficient simply to touch the tip of the tongue to the 
substance; unless some caution is given in the text, it may 
be considered safe to taste all substances studied this year. 
The odor is generally described by comparing it with some 
familiar odor from every-day experience; a substance that 
can be distinguished easily from all other substances by the 
odor alone is said to have a characteristic odor. By solu- 
bility is generally meant solubility in water; to be soluble 
a substance when put into water must lose its original form, 
disappear, and give a clear, tho not necessarily colorless, 
solution; heating a little often aids the dissolving; if the 
original substance is a powder it may form a cloudy or 
muddy mixture with water; if after standing a while this 
powder settles to the bottom it was not soluble in water. 
By specific gravity is meant the number of times heavier a 
definite volume of a substance is than an equal volume of 
some other substance taken as a standard; water is usually 
taken as a standard for solids and liquids, while gases are 
compared with air; the term density is often used instead 
of specific gravity, tho physicists distinguish carefully be- 
tween these two terms; in our work specific gravity can be 
determined only roughly, i.e., copper is lighter or heavier 
than some other metal with which you are familiar. The 
fusibility, or tendency to melt, should be tried first in the 
Bunsen burner flame, and then if necessary in the blast 
lamp flame. By inflammability is generally understood 
burning with a flame in ordinary air. The boiling of a sub- 
stance cannot be studied, of course, unless the substance 
can be melted or is already in the liquid state; the boiling 



irse, 



FIRST YEAR CHEMISTRY 



23 



point, or the exact temperature at which the substance boils, 
cannot always be determined easily; therefore, note sim- 
ply, if the substance boils, whether it boils easily or with 
difficulty. The volatility, like the fusibility, should be test- 
ed first in the Bunsen flame and then in the blast lamp; 
the substance is volatile if it disappears in vapor as water 
does when it goes off in steam; in order to call a substance 
volatile it must be possible to con- 
dense its vapor and get the origi- 
nal substance again. Consistency 
i> a property that is worth looking 
for in some substances; b}^ con- 
sistency is meant "the condition 
of standing or adhering together"; 
firmness and solidity are synony- 
mous with it; in regard to this 
property the substance in hand 
is generally compared with one 
of a list of substances, — as 
wax, putty, molasses, oil, and 
water. 

The blast lamp and its use. — Since the blast lamp was 
mentioned in the preceding paragraph a few words as to its 

use may not be out 
of place. The blast 
lamp is practically 
a Bunsen burner. 




Fig. 26. Blast lamp. 




=^SS3^^ 



~N » 

- E with a forced draft 



*s. 



of air to increase 
Fig. 27. Interior construction of the blast lamp. the rapidity of the 

combustion of the gases and, consequently, the heat 
of the flame. Fig. 26 shows a perspective view of the 
blast lamp, while Fig. 27 gives some idea of its interior 
construction. The rubber hose supplying the illuminating 



24 



FIRST YEAR CHEMISTRY 




Fig. 28. Foot bellows. 



gas is attached at A, and the one supplying compressed air 
at B. The supply of gas may be regulated either at the 

tap from which the gas 
is taken, or by the tap 
at C; the supply of com- 
pressed air is generally 
regulated by the tap D. 
The principle of the 
lamp is to force an 
extra amount of air into 
the center of the burn- 
ing gas. The air under 
pressure may be obtained 
by means of a foot bel- 
lows. See Fig. 28. By 
means of the leather bel- 
lows, air is pumped 
into the rubber bag, from which the air under pressure is 
conducted to the blast lamp by means of a rubber tube at- 
tached to the outlet. 

If the laboratory is supplied with water under pressure 
a much more convenient means of getting compressed air 
is the Richards blower, shown in Fig. 29. Screw the end A 
to the faucet and turn on the water. When the water rushes 
thru the aspirator B it draws air in at C. The wavy pipe 
D below the aspirator B, is to render the water more effec- 
tive in drawing air in thru C. The mixture of air and water 
passes into the large iron standpipe E; there they separate, 
the air going to the top, where it may be drawn off thru the 
outlet, regulated by the stop-cock G; the water goes to the 
bottom and escapes by the pipes H and / into the sink or 
waste drain. 

How to light the blast lamp. — Open the stop-cocks C and 
D in Fig. 26 and G in Fig. 29. Turn on the gas and light 



FIRST YEAR CHEMISTRY 



25 



The flame will be large, 
due to incomplete corn- 



it at the outlet E of the lamp. 
Luminous, and somewhat smoky, 
bustion of the gas. Turn on the 
water full at the faucet and wait 
a moment till an even pressure 
of air is obtained. If the flame 
is still luminous, turn off the sup- 
ply of gas till the yellowness just 
disappears from the flame. If a 
small tongue of yellow flame rises 
from the main flame just beyond 
the outlet E of the lamp, push 
forward the outside tube of the 
lamp, F in Fig. 26, till this tongue 
just disappears. For heating cru- 
cibles use the largest blue flame 
obtainable. If a very small flame 
is needed, turn off both air and 
water till the desired flame is ob- 
tained. The temperature obtain- 
able with an ordinary blast lamp 
varies from 1000° to 2000° Centi- 
grade. 



Experiment 6. Heating copper 
in contact with air. Have ready 
a porcelain crucible, pipe-stem 
triangle, tripod, Bunsen burner, 
horn-pan balance, set of smaller 
weights, and brass forceps; also 
some copper turnings. 

Fill a clean and dry porcelain crucible about two thirds 
full of copper turnings, taking care that the turnings are 
dry, clean, and entirely free from chips, dirt, and other 



A op 
d— f 

DM 

t 


F. 

||<3 


H 


II 



Fig. 29. Richards blower. 



26 



FIRST YEAR CHEMISTRY 



foreign matter. The cover of the crucible should not be 
used in this experiment. Get the exact weight of the cru- 
cible and turnings, using the horn-pan balance and set of 
smaller weights. Weigh accurately to centigrams, and re- 
cord the weights in the notebook as follows: 



Weight of crucible and copper after heating 
Weight of crucible and copper before heating 



g- 



Increase in weight 



g- 



Set the crucible on a pipe-stem triangle supported on the 

tripod, as shown in Fig. 30. Heat 
with the colorless Bunsen flame for 
about twenty minutes, beginning 
with a small flame, which just 
touches the bottom of the crucible, 
and increasing the size of the flame 
as the crucible gets hot enough to 
stand the high heat. Note any 
change in the appearance of the 
copper as it is being heated. Let 
the crucible cool till it can be 
handled with the fingers; then 
weigh it again and record the 
weights. How much was the in- 
crease in weight? This increase 
in weight shows that something 
came from the air and attached itself to the copper. It 
will be found that other metals act the same way when 
heated in the air. 

The substance that comes from the air and attaches it- 
self to the copper is called oxygen. In a later experiment 
oxygen will be made, and its properties studied in detail. 
What change in appearance was there to the copper during 




Fig. 30. 

Heating copper turnings. 



FIRST YEAR CHEMISTRY 27 

the heating? The black coating is a new substance and it 
contains copper and oxygen; it is called copper oxide or 
oxide of copper; the process of heating a metal and form- 
ing its oxide is called oxidation ; in this case copper is said 
to be oxidized when it is heated. 

There are two reasons for cooling the crucible before get- 
ting the exact weight after heating. What are they? Re- 
cord the answer in the notebook. 

Get the properties of copper oxide, comparing what you 
made with some copper oxide from the bottle on the shelf. 

Note on copper. — Copper is found in the ground in large quantities, 
both free and as minerals, in which it is combined with oxygen or 
sulphur. These minerals are called ores. The metallic copper may be 
obtained from the copper ores by heating to a high temperature with 
charcoal or coke; after that it is usually purified by means of electric- 
ity. The specific gravity of copper is about 8.5, its melting point is 
about 1100°C. and its boiling point is about 2100°C. Copper is used 
very widely in electrical work of all kinds, in boilers, in household 
utensils, in coins, and in such alloys as brass and bronze. 

Experiment 7. Zinc and its properties. Take zinc in 
its various forms, — sheet, mossy, and dust. Examine the 
different forms of zinc for the properties mentioned in the 
experiment on copper; refer to Experiment 5 if you do not 
remember all the properties or how to test for them. Re- 
eord all your results as usual in the notebook. If any of 
the properties vary with the different forms of zinc, be sure 
to record all such variations. 

Experiment 8. Heating zinc in contact with air. Have 

ready a Hessian crucible, ring stand with small ring, iron 
rod. and blast lamp; also some mossy zinc. 

Put enough mossy zinc in a small Hessian crucible to 
half fill it, support it in the small ring of the ring stand, 
and heat it over the blast lamp till the zinc melts. Continue 



28 FIRST YEAR CHEMISTRY 

the heating, with an occasional stir by means of an iron rod, 
till the zinc burns with a flame. What is the color of the 
flame? What happened chemically to the zinc when it 
burned? What, then, is the burning, chemically? As a 
help in answering these two questions, refer to what hap- 
pened to copper when it was heated. Where did the oxy- 
gen come from that united with the zinc during its burning? 
What name would you give to the powder left after the 
burning? Examine the product of combustion, which some- 
times forms what is called "Philosopher's wool," a light, 
fluffy form of zinc oxide. Get the properties of zinc oxide, espe- 
cially its color when hot and when cold, comparing that which 
you made with the zinc oxide from the bottle on the shelf. 

Note on zinc. — Unlike copper, zinc is never found free in the ground. 
Its common ore is zinc blend, a compound of zinc with sulphur; an- 
other ore is composed of zinc, carbon and oxygen. The metallic zinc 
is obtained by roasting the ores in the air, whereby the zinc is changed 
to zinc oxide; this is then heated with charcoal in fireclay crucibles, 
and the zinc vapor formed passes over into iron receivers where it con- 
denses first as zinc dust and finally as fused, compact metallic zinc. 
The specific gravity of zinc is about 7; it melts at about 420°C; and 
boils at about 920°C. Zinc is used extensively in electric batteries, in 
galvanizing iron, in making brass and bronze, in housebuilding, and 
for household purposes. 

Experiment 9. Magnesium and its properties. Take mag- 
nesium in its two forms, — ribbon and powder. The former 
may be obtained from the instructor, and the latter may be 
found in bottles on the shelf. Examine the two forms of 
magnesium for all the properties mentioned in Experi- 
ment 5. Note particularly the color, luster, tenacity, brit- 
tleness and density. Do any of the properties, such as 
luster, vary between the two forms of magnesium? 

Experiment 10. Heating magnesium in contact with air. 
Have ready a Bunsen burner, horn-pan balance, set of smaller 



FIRST YEAR CHEMISTRY 29 

weights, forceps, porcelain crucible, pipe-stem triangle, 
tripod, and iron rod; also some magnesium ribbon and 
powder. 

Hold a short piece of magnesium ribbon in the Bunsen 
flame until it bursts into flame itself. What is the color of 
the residue? What is the chemical name of the residue? 
What happened to the magnesium chemically when it 
it burned? What two marked differences are there between 
magnesium and copper in the formation of the oxide? What 
similarity is there chemically between magnesium and cop- 
per in the formation of the oxide? What similarity is there 
between magnesium and zinc in the formation of the ox- 
ides? Examine the magnesium oxide and find all its prop- 
erties you can; compare that which you made with the 
magnesium oxide from the bottle on the shelf. 

Fill a clean and dry porcelain crucible about one third 
full of magnesium powder. The cover of the crucible should 
not be used in this experiment. Get the exact weight of 
the crucible and powder, weighing accurately to centi- 
grams, and recording the weights in the notebook as fol- 
lows: — 

Weight of crucible and magnesium after heating = g. 

Weight of crucible and magnesium before heating = g. 



Increase in weight = g. 

Support and heat the crucible as explained in Experiment 6, 
where copper was heated in air. Heat at least ten min- 
utes but do not stir the contents of the crucible; then cool 
and weigh. How much was the increase in weight? When 
copper was heated what caused the increase in weight? 
What, then, probably caused the increase in weight when 
magnesium was heated? Does the oxide of magnesium 
made from the powder differ from that made from the 



30 FIRST YEAR CHEMISTRY 

ribbon? Record any difference in properties between the 
two oxides of magnesium you made. 

Note on magnesium. — Magnesium is never found free in the ground, 
but it occurs in combination with carbon and oxygen in a rock resem- 
bling marble. When combined with sand magnesium occurs in such 
rocks as talc, soapstone, and asbestos. Magnesium often accompanies 
lime in its compounds. Another source of magnesium is a soluble 
compound of magnesium and potash found largely in Germany; from 
this compound the metallic magnesium is obtained by means of elec- 
tricity. The specific gravity of magnesium is 1.75; it melts at about 
700°C. and boils at about 1100°C. Its principle use is in flash-light 
powder and in fireworks. 

Experiment 11. Phosphorus and its properties. Read 
the directions for this experiment way thru before starting 
to do any of the work. Phosphorus occurs in two forms, — 
red and yellow. Both forms may be found in the hood, 
and all work with phosphorus must be done in the hood instead 
of at the individual desks. 

Caution. — Extreme care must be taken when working 
with phosphorus. The yellow variety is very inflammable; 
for this reason it must be kept under water when not in 
use; also it must be cut under water. Burning phosphorus 
sticks to the skin tenaciously, and the burn left is generally 
more troublesome than an ordinary burn. Yellow phos- 
phorus must always be handled with the forceps, never 
with the bare fingers; the red phosphorus, however, may 
safely be handled with the fingers. 

Have ready a Bunsen burner, test tube rack, two large 
test tubes, graduate, forceps, knife, filter paper, and the 
hood; also phosphorus, both red and yellow. 

Take out a little red phosphorus and examine it for the 
properties mentioned in Experiment 5. It is best in this 
experiment to arrange the properties in list form so that 
the properties of yellow phosphorus may be put alongside 



FIRST YEAR CHEMISTRY 



31 



those of the red for comparison. Examine the yellow 
phosphorus while it is still under water, and note as many 
of its properties as you can. Then take out a small piece 
from the water, and get as many more properties as you 
can; handle it with the forceps, never with the bare fingers. 
The phosphorus may be dried by pressing it gently between 
filter paper. In the case of yellow phosphorus look par- 
ticularly for the consistency, because this substance is the 
first good illustration we have had of this property; in the 
case of red phosphorus, the consistency is practically iden- 
tical with the existence in the powder state. 

Melting point of phosphorus. — That phosphorus melts 
may best be shown as follows: Put a piece of yellow phos- 
phorus about the size of 
a pea in the bottom of a 
large test tube and add 
10 c.c. of water. Make 
a test tube holder by 
folding a piece of paper 
several times lengthwise, 
putting it around the test 
tube next to the flange 
at the top of the tube, 
and pinching it tightly 
between the thumb and 
fingers as shown in Fig. 
31. If the test tube hap- 
pens to be wet on the 
outside, wipe it dry. 
Heat the water gently, 
hold the test tube in the 

colorless Bunsen flame, 91 TT . + , . , , ., 

, Fig. 31. Using the test tube holder, 

first passing the test 

tube back and forth a few times in and out of the flame 



1 




^ 



32 



FIRST YEAR CHEMISTRY 



till the test tube has become a little warm. Watch 
the piece of phosphorus closely for any change in state, 
and note about how much time elapses from beginning the 
heating till the phosphorus melts. When the phosphorus 
has melted, set the tube aside in the rack. Then heat 10 
c.c. of water in another test tube, this time without phos- 
phorus, and note about how long it takes to bring the water 
alone to a boil. The melting point of ice is 0°C, the boil- 
ing point of water is 100°C, and ordinary temperature 
varies from 15° to 20°C. Estimate roughly from the times 
noted above at about what temperature you think phos- 
phorus melts. 



Experiment 12. Heating phosphorus in contact with air. 

Have ready two dry fruit jars of one pint size and of " Light- 
ning" make, Bunsen burner, deflagrating spoon, horn-pan 
balance, set of smaller weights, forceps, knife, filter paper, 
porcelain crucible, graduate, pneumatic trough, and the 
hood; also some phosphorus, both red and yellow, and 
some vaseline. 

Do this experiment in the hood. 

See that the rubber washers are new or in good condition; 
if you have any doubt about their ability to render the jars 
air-tight, take a little vaseline between the thumb and fore- 
finger and moisten the rubber with it, but do not have any 
lumps of vaseline remaining on the rubber. Also see that 
the deflagrating spoon is clean and dry. 

Either kind of phosphorus is suitable for this experiment, 
but the yellow variety is more satisfactory to work with. 
If you use the red variety, weigh out just a half a gram of 
it in the horn-pan balance. If you use the yellow variety, 
put the 5 decigram weight on the right-hand pan of the bal- 
ance; then with the forceps take out a piece of phosphorus 
which you think is of about the right weight, dry it a little 



FIRST YEAR CHEMISTRY 33 

by pressing it gently between filter paper, and, still hand- 
ling it with the forceps, put it in the left-hand pan of the 
balance. If the piece is too heavy, cut it under water to 
the right size, dry it, and weigh it again. When you have 
trimmed it to weigh half a gram, note the size of it, because 
future experiments will need the same amount of phosphor- 
us, and in those cases it will be sufficient to estimate the 
amount needed if you remember the size of a half gram 
piece. 

When the phosphorus is all ready, put it in the deflagrat- 
ing spoon; have ready a lighted Bunsen burner over which 
the spoon may be held to light the phosphorus. Hold the 
spoon in one hand and the cover to the jar in the other. 
Light the phosphorus, and quickly but carefully lower the 
spoon into the jar; at once put on the cover and snap it 
down. Tho there is no great danger of the jar breaking, 
it is just as well to step back a little; if the jar was not dry 
inside, the drops of water in contact with the hot glass might 
crack the jar; the same might happen if the melted phos- 
phorus should happen to run off the spoon when it is low- 
ered into the jar; in case the jar cracks and the phosphorus 
comes out, the burning phosphorus is likely to spatter 
around, but if the jar stands on the soapstone bottom of 
the hood and the window of the hood is shut down there is 
no danger. 

Note, as the phosphorus burns away, the white powder 
formed. When copper was heated in air, what was the 
black substance formed? When magnesium burned in air, 
what was the white residue? What, then, would you call 
the white powder formed in this experiment? What hap- 
pened chemically to the phosphorus when it burned? Where 
did the phosphorus get the oxygen that it united with? 

When the phosphorus has stopped burning and the jar 
has cooled, pry the cover off the jar; then get all the prop- 



34 



FIRST YEAR CHEMISTRY 



erties you can of the phosphorus oxide. What caused the vacu- 
um in the jar ? Compare the phosphorus oxide you made with 
the sample of phosphorus oxide in the bottle on the shelf. 

Clean the deflagrating spoon from any phosphorus remain- 
ing upon it by holding it in the Bunsen flame till all the 
phosphorus has burned off. Every time you use phosphorus 
be sure to clean the spoon this way before putting it away. 
Proportion by volume of oxygen in air. — Thus far we have 
found four substances that take oxygen from the air when 
they are heated or burned. Of these, phosphorus unites 
with oxygen the most vigorously; it can, therefore, be 
used to find out what proportion by volume of the air is 
oxygen. To determine this, proceed as follows: Fill the 
pneumatic trough to the brim with water and float an empty 
crucible on the surface of the water. Select a piece of yel- 
low phosphorus weighing about half a gram and put it in 
the bottom of the crucible. Hold a pint fruit jar inverted 

full of air but with 
the cover off in 
one hand, while 
with the other 
you light the 
phosphorus by 
touching alighted 
match to it. As 
soon as the phos- 
phorus is burning 
well, bring the 
inverted jar 
down over the 
crucible, dip the 




Fig. 32. 
Determining the proportion of oxygen in air. 



mouth of the jar just under the surface of the water, 
as shown in Fig. 32, and hold it there till the phos- 
phorus has stopped burning. What are the white fumes 



FIRST YEAR CHEMISTRY 35 

that fill the jar? Does the water rise any in the jar? 
Reach in the fingers, tip over the crucible, and let it sink to 
the bottom of the trough. Raise or lower the jar as neces- 
sary till the level of the water is the same inside the jar as 
outside; snap on the cover and remove the jar from the 
water. Shake well till the white fumes have disappeared. 
What does this show regarding their solubility in water? 
Open the jar, measure with the graduate the water that ran 
into the jar, and record its volume in the notebook. Fill 
the jar with water to the brim and measure the volume of 
this water. Find out, by dividing one figure by the other, 
about what part of the air in the jar was occupied by the 
oxygen. Was part of the air left in the jar because not 
enough phosphorus was used? To answer this, repeat the 
experiment, using about three times as much phosphorus, 
and see if three times as much water runs in. What pro- 
portion by volume, then, of oxygen is there in the air? 
Note that the gas left in the jar after dissolving the fumes 
of phosphorus oxide is colorless. A study of this part of 
the air will be made as soon as we have finished the work 
with phosphorus. 

General considerations. — Having studied the four sub- 
stances, copper, zinc, magnesium, and phosphorus, together 
with the oxidation of these four substances in air, we are 
now in a position to consider several general statements and 
definitions. These are contained in the following paragraphs. 

A simple substance is one that cannot be broken down into 
simpler substances. For instance, a piece of copper, no 
matter how finely it may be divided, can never be divided 
into particles containing anything but copper. For that 
reason it is called a simple substance. In like manner zinc, 
magnesium, and phosphorus are called simple substances. 
A >imple substance is sometimes spoken of as an elemen- 
tary substance, or simply an element. 



36 FIRST YEAR CHEMISTRY 

A compound substance is one that contains two or more 
simple substances united in such a way that the compound 
substance has properties that are different from those of either 
of the simple substances of which it is composed. For instance, 
copper oxide is a compound substance, consisting of copper 
and oxygen. As discovered in Experiment 6, it is a black 
powder, whereas the copper itself is a reddish brown solid, 
and the oxygen coming from the air is a colorless gas. A 
compound substance may often be made by the direct 
union of the simple substances; in copper oxide the oxygen 
is not present in its elementary state; its presence may be 
shown, however, by liberating it by proper means from the 
oxide. Later experiments will deal with such decomposi- 
tion of compound substances. Of course,, all the oxides 
you have made are compound substances. 

When a new substance is taken up for study, the first 
question naturally is: "Is the substance simple or com- 
pound"? All metals are simple substances. Some non- 
metallic substances, such as phosphorus, are also simple. 
Most of the substances made or used in the laboratory are 
compound substances; in such cases the first question is: 
"What simple substances are there in the compound"? 
The composition of the compound may be proved either by 
building it up from the simple substances or by taking 
some of the compound and breaking it down into the simple 
substances of which it is composed. 

Physical changes and chemical changes. — When a sub- 
stance melts, as phosphorus did when it was heated under 
water, and as zinc did when it was heated in the Hessian 
crucible, there is a change in the state or form of the sub- 
stance, but not in the composition of it. The phosphorus 
was just as much phosphorus in the liquid state as in the 
solid; so was the zinc. Such a change is called a physical 
change, because it affects only the physical properties of 



FIRST YEAR CHEMISTRY 37 

the substance. Any change, then, which does not affect 
the composition of the substance is called a physical change. 
On the other hand, when phosphorus burned, its oxide 
was formed, i.e., there was a change from the simple sub- 
stance phosphorus to the compound substance phosphorus 
oxide, containing the two simple substances, phosphorus and 
oxygen. Such a change is called a chemical change, be- 
cause it changes the composition of the substance started with. 
Any change, then, which does affect the composition of the sub- 
stance is called a chemical change. Most of the work for this 
year will have to do with chemical changes; these, however, 
are often accompanied or indicated by physical changes. 

Equations. — It is convenient to have some short method 
for recording a chemical change to avoid the necessity of 
explaining each change in full. For this purpose an equa- 
tion is used. An equation is, then, a short-hand expres- 
sion of the changes in composition brought about during 
a chemical change. Thus, the fact that when copper is 
heated, oxygen from the air unites with it, forming copper 
oxide, is expressed as follows: 



copper [ + j oxygen } = { ™PP^ } 



and is read " copper plus oxygen equals copper oxygen." 
Substances started with are always put on the left-hand 
side of the equation and the substances obtained are always 
put on the right. In the above equation the brackets 
around copper indicate that the simple substance, metallic 
copper, was started with; in like manner, the brackets 
around oxygen show that this substance was used in its 
simple gaseous state. The fact that the compound formed 
contains the two simple substances, copper and oxygen, is 
indicated by putting the names of these two substances 
one over the other within brackets. The name of the com- 



38 FIRST YEAR CHEMISTRY 

pound substance, in this case, " copper oxide," is never put 
inside brackets. Every simple substance appearing on one 
side of the equation must appear on the other side also; 
in this respect a chemical equation is like an algebraic equa- 
tion, i.e., it must balance; in all other respects it differs 
from an algebraic equation. 

Factors and products. — A factor is any substance started 
with, and a product is any substance formed in a chemical 
change. In the equation in the preceding paragraph there 
are two factors, copper and oxygen, and one product, 
copper oxide. In this same equation, the two factors are 
both simple substances, while the product is a compound 
substance. The number of products in any chemical change 
depends on the factors used, and the products are not nec- 
essarily always compound substances. 

Write equations similar to the one above, for the oxidation 
of zinc, for the burning of magnesium ribbon, and for the 
burning of phosphorus. In all future experiments wher- 
ever there is a chemical change try to write the equation, 
remembering always that the equation is simply a short 
method for expressing what you have learned by experi- 
ment about the substances, and that the equation in it- 
self never proves anything regarding the composition of a 
substance. 

Note on phosphorus. — Phosphorus never occurs free in nature; 
many phosphates- — compounds composed of metals, phosphorus and 
oxygen — are found in the ground, phosphate of lime being the most 
common one; the bones of animals consist chiefly of phosphate of lime. 
Phosphorus may be made by burning old bones, treating the residue 
with sulphuric acid, filtering, mixing the filtrate with charcoal, evap- 
orating to dryness, and then heating to a high temperature in clay 
retorts; the phosphorus distills over as a vapor, which is condensed 
under water as a yellow solid; phosphorus may also be made in the 
electric furnace. Red phosphorus is made by heating the yellow va- 
riety in air-tight retorts for some time at about 260°C. The specific 



FIRST YEAR CHEMISTRY 39 

gravity of yellow phosphorus is about 1.8, of the red about 2.2; yellow 
phosphorus melts at about 44°C. and boils at about 290°C; red phos- 
phorus does not melt. Phosphorus is used chiefly in making matches. 
Air and its composition. — By air is meant the layer of gases cover- 
ing the surface of the whole earth and extending out into space some 
fifty miles or more. The air taken as a whole is sometimes called 
the atmosphere. The composition of air is practically constant, being 
approximately 78 parte of nitrogen, 21 parts of oxygen, and 1 part of 
argon. Several other substances, such as water vapor, carbon diox- 
ide, ammonia and dust, are often present, but in such small quanti- 
ties that they may be considered impurities and neglected in our con- 
sideration of the subject. For a long time it was thought that the 
air consisted simply of about one fifth oxygen and four fifths nitrogen, 
but in 1894 Rayleigh and Ramsay, two English scientists, noticed 
that nitrogen obtained from the air was a little heavier than nitrogen 
made from nitrogen compounds. Further work with atmospheric 
nitrogen yielded for them a small amount of a hitherto unknown 
gas, which they called argon. By passing clean, dry air over hot, 
finely divided copper the oxygen is removed; by passing the residue 
over heated magnesium powder the nitrogen is removed and the ar- 
gon is left. All three gases, oxygen, nitrogen, and argon, are color- 
less; oxygen is a good supporter of combustion; the others are in- 
ert gases. Liquid air may be made by subjecting clean, dry air to 
tremendous pressure and low temperature. It is a pale blue liquid, 
resembles water, is slightly heavier than water, and boils at about 
— 190°C. under atmospheric pressure. On exposure to ordinary air 
it evaporates rapidly and completely. Its principal uses are the prep- 
aration of pure oxygen and in the study of chemical changes at ex- 
tremely low temperatures. Air is really a mixture of gases, altho we 
speak of its ''composition"; the main reasons for considering air a 
mixture and not a definite chemical compound are: (1) The propor- 
tion of oxygen, of nitrogen and of argon is not absolutely uniform, 
but varies between small limits, these differences, however, being 
greater than those found in different analyses of a pure chemical com- 
pound; (2) If the various components of the air are mixed in the 
proportions in which they are found in the atmosphere, the result is 
exactly like air but there is no evidence of reaction, such as libera- 
tion of heat or production of light; (3) A greater proportion of oxy- 
gen than of nitrogen dissolves, when air is dissolved in water, whereas 
if these two gases were chemically combined in the air, the dissolved 
air would, of course, have the same composition as air itself; and 



40 FIRST YEAR CHEMISTRY 

(4) When liquid air that has been made from atmospheric air is al- 
lowed to evaporate, the nitrogen evaporates first, leaving the oxygen, 
whereas if the air were a chemical compound, it would have a definite 
boiling point at which it would evaporate unchanged in composition. 

Experiment 13. Iron and its properties. Take iron in 
its various forms, — sheet, filings, wire and "steel wool." 
Examine the different forms of iron for the properties men- 
tioned in the experiment on copper. 

Experiment 14. Heating iron in contact with air. Have 
ready a porcelain crucible, pipe-stem triangle, tripod, Bun- 
sen burner, horn-pan balance, set of smaller weights, and 
forceps; also some iron filings. 

Fill a clean and dry porcelain crucible about two thirds 
full of clean iron filings, making sure that they are free 
from chips, dirt, and other foreign matter; see also that 
they have not already started to rust. The cover of the 
crucible should not be used in this experiment. Get the 
exact weight of the crucible and filings, weighing accu- 
rately to centigrams, and recording the weights in the note- 
book as follows: — 

Weight of crucible and iron filings after heating = g. 

Weight of crucible and iron filings before heating = g. 



Increase in weight = g. 

Set the crucible on the pipe-stem triangle, supported on the 
tripod, just as in the experiments on heating copper and 
magnesium. Heat with the Bunsen burner for about 
twenty minutes, let the crucible cool till it can be handled 
with the fingers, and then weigh it again. How much was 
the increase in weight? What caused this increase in weight? 
Where did the oxygen come from? What change in ap- 
pearance was there to the iron as a result of the heating? 



FIRST YEAR CHEMISTRY 41 

What is the chemical name of the compound formed? Ex- 
amine the iron oxide and get all the properties you can of 
it; compare the iron oxide you made with the two samples 
of iron oxide from the bottles on the shelf. 

Write the equation that represents the change that takes 
place when iron is heated in air. 

Note on iron. — A little free iron occurs in nature, the most of this 
kind, however, being found in meteorites. Iron ores are numerous 
and important; most of them are oxides of iron of varying composi- 
tion and are called hematite, limonite, and magnetite; another ore 
is called siderite and contains iron, carbon, and oxygen. Still another 
common ore is pyrites — composed of iron and sulphur — but this one 
is not well adapted for iron making. Metallic iron is obtained from 
its oxygen ores by roasting them in air and then heating them with 
coke in furnaces which allow the melted iron to run off at the bottom. 
Ordinary iron is never pure; the two principal impurities are carbon 
and manganese, small quantities of sand, sulphur, and phosphorus 
being sometimes present. The different kinds of iron depend upon 
the percentage of carbon present; cast iron contains from 1.5% to 
6.0% of carbon, wrought iron from none to 0.5%, and steel from 0.5% 
to 1.5%. The specific gravity of iron varies as follows: cast iron 
from 7.0 to 7.6, wrought iron from 7.25 to 7.78, and steel from 7.6 
to 7.8; cast iron melts at about 1050°C, wrought iron at about 1500°C, 
and steel at about 1300°C; the boiling point of iron has not been de- 
termined. Iron is very extensively used in the construction of large 
buildings, in railroading, in ships, in machinery, and in many smaller 
articles, too numerous to mention. Iron forms two series of com- 
pounds, — the ferrous and the ferric compounds. More will be said 
about this later when the compounds of iron are taken up in detail; 
the extraction and working of iron will be studied at greater length 
at that time. 



Experiment 15. Mercury and its properties. Read the 
directions for this experiment way thru before starting to 
do any of the work. 

Caution. — Mercury in its ordinary state is not a poison; 
it may be handled with the fingers without danger. (The 



42 FIRST YEAR CHEMISTRY 

statement that " mercury is a poison" generally refers to 
the action on the human system of soluble compounds of 
mercury when taken internally.) Care should be taken, 
however, when mercury is heated, not to inhale any large 
amount of the vapor, as this has a poisonous action which 
affects the mucous membrane of the nose and lungs. 

Mercury may be found in bottles on the shelf. Take a 
small globule of it and examine it for its properties. What 
three properties of mercury are most prominent? Compare 
it with each of the simple substances already studied and 
see which properties are more marked in the case of mercury. 

To test the boiling and volatility of mercury put a glob- 
ule the size of a pea in a dry test tube, and, using the test 
tube holder described in the experiment on phosphorus, 
heat gently till the mercury boils; move the test tube con- 
stantly in the flame to keep the glass from breaking. Note 
the deposit on the sides of the glass in the upper part of 
the tube. What is it? Rub a glass rod over the deposit 
to collect it, if there is not much of it. Has the vapor of 
mercury any color? In what part of the test tube did the 
vapor of mercury exist? Does mercury boil at a higher or 
lower temperature than water? To find this out, heat in 
another test tube a drop of water as large as the mercury 
globule used and note about how soon that boils. Does 
the mercury oxidize, as far as you can see, when it is heated ? 
What was there about this experiment which showed that 
mercury does not oxidize when it is heated, as copper and 
zinc do when they are heated? 

Surface tension. — Mercury exhibits, better than any other substance 
we study this year, the property called surface tension. By surface 
tension we mean practically surface tenacity. The globule of mercury 
may be likened to a toy balloon; in the place of the thin and tena- 
cious film of rubber enclosing the spherical mass of air with which it 
is blown up we have a spherical globule of mercury surrounded by a 



FIRST YEAR CHEMISTRY 43 

film of mercury, in which film the liquidity of the metal is lost sight 
of in the presence of the highly developed tenacity thruout this film 
of mercury. The influence of the surface tension keeps the mercury 
always in spherical form unless the mass of mercury is so large that 
its great weight overcomes the surface tension and makes the mer- 
cury take on the shape of the containing vessel. All liquids show 
some surface tension when in small drops. 

Experiment 16. Mercury oxide. The preceding experi- 
ment showed that mercury boils easily without oxidizing, 
i.e., that mercury oxide is not formed when mercury is 
heated in a test tube. This might indicate that mercury 
oxide is not known. Mercury oxide is known, however. 
Prolonged heating of mercury near its boiling point will 
produce a little oxide of mercury. The oxide is more con- 
veniently made on a larger scale by decomposing compli- 
cated mercury compounds, and the oxide so made finds 
considerable use in the laboratory. 

Take a very little mercury oxide from the bottle on the 
shelf and get as many of its properties as you can. Do not 
heat it, however, for the effect of heat on it is to be studied 
in detail in the next experiment. 

Experiment 17. Preparation of oxygen from mercury 
oxide. Have ready a Bunsen burner, triangular file, some 
short pieces of glass tubing, and several wooden toothpicks; 
also some mercury oxide. 

How to make a bulb tube. — First make a bulb tube as 
follows: Cut off a piece of glass tubing about 7 cm. long. 
Hold one end of it in the hottest part of the Bunsen flame, 
turning it continually between the fingers. When the end 
of the tube has softened enough to close the end of the 
tube entirely, remove it from the flame and blow in the other 
end. The melted glass swells out to form a bulb; this 
bulb should have a diameter not more than twice as great 



44 



FIRST YEAR CHEMISTRY 



Fig. 33. 
Glass tube with end fall- 
en together just a little. 



Fig. 34. 
Glass tube just ready to 



close. 



Fig. 35. 
Finished bulb tube. 



as the diameter of the tube; if the bulb is much larger 

than this, the walls are so thin that 

* they fall together when the tube is 

reheated. Fig. 33 shows a tube the 
sides of which have fallen together 
just a little; Fig. 34 shows the same 
tube when the opening is just ready 
to close; and Fig. 35 shows the bulb 
tube complete. If the inside of the 
bulb is not perfectly dry, roll up a 
narrow strip of filter paper and dry 
out the tube with that. Put a little 
mercury oxide in the bulb tube, 
enough to half fill the bulb, but do 
not have any in the stem. If any of 
the mercury oxide has stuck to* the 

sides of the tube, roll up a piece of filter paper and wipe out the 

inside of the straight part of the bulb tube. Heat the tube 

gently in the Bunsen flame, and note whether there is any 

change in color to 

the mercury ox- 
ide. Continue 

the heating for 

some time and 

occasionally 

plunge into the 

open end of the 

bulb a wooden 

toothpick that is 

glowing at the 

end, as shown in 

Fig. 36. Do not FlQ 36 Testing a gas with a g i ow i ng splinter. 

drop the tooth- 
pick into the tube; simply hold it at the open end of the tube. 




FIRST YEAR CHEMISTRY 45 

Note what happens to the glowing splinter. Has any deposit 
collected on the sides of the bulb tube? If there has, what 
is it? Break the tube, if necessary, so as to scrape the de- 
posit together. When that portion of the mercury oxide 
left in the bulb unchanged by the heat has cooled, examine 
it and see if it has regained its original color. 

The breaking down of mercury oxide into mercury and 
oxygen may be represented by an equation, the substance 
started with being put on the left-hand side of the equality 
sign, and the substances obtained being put on the right- 
hand side, thus: 

( mercurv ) ( I i J ! 

' J \ = 4 mercurv - -j- < oxygen > 

( oxygen j ( J j ' ( J& j 

The decomposition of a compound substance into the simple 
substances of which it is composed is called analysis; it is, 
in a way, the opposite of oxidation. 

Note on mercury. — The common name for mercury is quicksilver. 
Only occasionally is free mercury found in rocks. The most com- 
mon ore is cinnabar, a compound of mercury with sulphur. Metallic 
mercury is made by roasting cinnabar, whereby the sulphur burns 
off, and the mercury vapors are condensed in large chambers. It is 
often purified by distilling or by filtering thru chamois skin. The 
specific gravity of mercury is 13.6; solid mercury melts at about — 39°C; 
liquid mercury boils at about 360°C. The main uses of mercury are 
in making thermometers and barometers, and in extracting gold and 
silver from their ores. 

Experiment 18. Preparation of oxygen from potassium 
chlorate. Have ready a Bunsen burner, test tube, test tube 
rack, and several wooden toothpicks; also some well crystal- 
lized potassium chlorate. 

Potassium chlorate is a compound substance containing 
three simple substances, — potassium, chlorine, and oxy- 
gen. When it 8 heated, it decomposes readily, the oxygen 



46 FIRST YEAR CHEMISTRY 

being set free in its gaseous state while the other two simple 
substances are left behind. 

Put a few crystals of potassium chlorate in a dry test 
tube, and heat in the Bunsen burner gently at first, and 
then with the full heat of the Bunsen burner. When the 
crystals have melted to a colorless liquid, continue the 
heating until white fumes begin to rise in the test tube, 
then plunge into the mouth of the test tube a wooden tooth- 
pick that is glowing at the end. What happens to the 
glowing splinter? What did this phenomenon indicate in 
the preceding experiment? What does it indicate in this 
experiment? What then, is the test for oxygen? Set the 
test tube aside in the test tube rack to cool. The white 
solid left in the test tube is called potassium chloride, and 
is composed of potassium and chlorine. The white fumes 
that arose in the test tube towards the end of the heating 
were largely potassium chloride. Since this substance vapor- 
izes easily, would you say that potassium chlorate vaporizes ? 
On what experimental fact do you base your decision? 

Write the equation for the decomposition of potassium 
chlorate by means of heat. 

Experiment 19. Preparation of oxygen from potassium 
ch 1 orate on a large scale. Have ready a Kjeldahl flask, 
with cork to fit it, ring stand, clamp, medium ring, Bunsen 
burner, four " Lightning " fruit jars with washers, pneumatic 
trough, large funnel, filter paper, glass rod, two large beakers, 
porcelain evaporating dish, horn-pan balance, set of smaller 
weights, forceps, tripod, gauze, wash-bottle, mortar and 
pestle, and some glass tubing of at least 5 mm. bore; also some 
potassium chlorate and some black oxide of manganese. 

Set up the apparatus as follows: Clamp the Kjeldahl 
flask to the ring stand at an angle of about 45° at such a 
hight that the body of the flask may be heated convenient- 



FIRST YEAR CHEMISTRY 



47 



ly by a Bunsen burner. Fit the flask with a good sound 
cork, thru which you have made a hole for the glass tube. 
Fill the pneumatic trough about two thirds full of water 
and set it at one side of the ring stand as shown in Fig. 37; 




Fig. 37. Apparatus for generating oxygen on a large scale. 

see that the bridge, the strip of sheet metal with a hole in 
the center of it, is in place. From the cork let a glass tube, 
starting flush with the inside of the cork, pass down into 
the pneumatic trough. The tube should be bent, as shown 
in Fig. 37, and the lower end of the glass tube should be 
bent so that the end just sticks up through the hole in the 
bridge. Fit the four fruit jars with tight washers. Fill 
each of the four jars completely with water, and putting 
the hand tightly over the mouth of each jar, invert them 
in the sink; be careful that no bubbles of air stay in the 
jars. When the apparatus has all been set up as shown 
in Fig. 37, proceed with the weighing of the chemicals. 



48 



FIRST YEAR CHEMISTRY 



Preparation of the chemicals. — Weigh out on the horn- 
pan balance 15 grams of potassium chlorate. Empty the 
chlorate into the dry porcelain mortar, and if the crystals 
are large crush them a little with the pestle. Next, weigh 

out exactly three grams of 
black oxide of manganese, 
weighing it carefully to a 
tenth of a gram. Then 
empty the black powder 
upon the potassium chlo- 
rate in the mortar, and 
mix the two as intimately 
as possible, not by grind- 
ing with the pestle, but 
by stirring with a glass 
rod. Put the mixture 
into the Kjeldahl flask, 
which should be perfectly 
dry; if any water has 




Fig. 38. 
Transferring a dry powder to a flask. 



got into the flask during the previous work with it, order 
another one from the storeroom, making sure that you get 
one that fits the cork you have prepared. Set the wet flask 
away to dry, — on the radiator, if steam is on the building, 
otherwise in any warm place you can find. The best way 
to empty a dry powder into a narrow mouthed flask is to 
transfer the powder to a sheet of white, glazed paper (filter 
paper will not do, because powders stick to its rough sur- 
face so easily), and then slide the powder from the paper 
directly into the flask as shown in Fig. 38. Connect the 
apparatus again as shown in Fig. 37, and set one of the 
inverted fruit jars filled with water on the pneumatic bridge 
over the end of the glass delivery tube. 

Generation of the gas. — Heat the Kjeldahl flask with 
the Bunsen burner, passing the flame lightly back and forth 






FIRST YEAR CHEMISTRY 49 

until the flask gets hot. Finally set the Bunsen burner 
underneath the flask and regulate the supply of gas to the 
burner so that the oxygen shall not be evolved too rapidly. 
When the first jar is nearly filled with gas, remove it from 
the bridge and snap on the cover, while the mouth of the 
jar is still under the surface of the water. Not more than 
a few c.c. of water should be left in the jar; the less water 
is left in each jar the better. Put another jar on the bridge, 
fill that with oxygen in the same way and seal it. Con- 
tinue until all four jars have been filled with oxygen or 
until no more gas is evolved. If no oxygen is wasted, at 
least four jarfuls ought to be obtained from the 15 grams 
of potassium chlorate used. When the oxygen has ceased 
coming off, loosen the cork from the Kjeldahl flask and 
then take away the Bunsen burner. It is important to 
loosen the cork before the flame has been taken away and 
not afterwards. Why is this so? Set the jars of oxygen 
to one side while you take care of the flask and its residue. 

Treatment of the residue in the flask. — Do not throw away 
the contents of the Kjeldahl flask, but when it has cooled 
a little add about 100 c.c. of warm water. The white chlo- 
ride of potassium dissolves in the water while the black 
oxide of manganese does not. The mixture should now 
be filtered in order to separate the two substances from 
each other. 

How to filter. — The method by which a liquid is sepa- 
rated from any solid substance with which it is mechani- 
cally mixed is that of filtering the mixture thru porous 
paper known as filter paper. This process is described in 
detail in the next three paragraphs. 

Fold a circular filter along the diameter AB in Fig. 39 
and get the doubled paper shown in Fig. 40. Fold the 
doubled paper along the line CD of Fig. 41, and got the 
form shown in Fig. 42. Open this in such a manner that 



50 



FIRST YEAR CHEMISTRY 



three thicknesses of paper form one half of the cone, while 
the fourth thickness of paper forms the other half as shown 
43. Insert the tip of the cone in the point of the 



in Fig. 




First step in folding 
a filter. 




B 



C 

Fig. 40. 

Second step in fold 

ing a filter. 




B 



C 

Fig. 41. 

Third step in fold 

a filter. 



funnel and fit the paper closely to the sides of the funnel; 
in order to make the paper stay in place, wet it slightly 
by means of the stream from the wash-bottle and press the 
paper tightly against the glass. Be careful to squeeze the 



C AB 

Fig. 42. 
Fourth step in 
folding a filter. 




c 

Fig. 43. 
A finished filter. 



folds of the paper flat against the glass. No bubbles of 
air should remain between the paper and the glass. If the 
funnel is of such an angle that it is impossible to fit a care- 
fully folded paper to it snugly, exchange the funnel at the 
storeroom for one having the proper angle. The paper 
must never stick up above the top of the funnel; and the 
upper edge of the paper ought to be from one to two centi- 
meters below the top edge of the funnel. 

Set the funnel to which the filter paper has been fitted 
in a ring of the ring stand as shown in Fig. 44; let the end 



FIRST YEAR CHEMISTRY 



51 



of the stem of the funnel touch the side of an empty beaker 
underneath. With a glass rod stir up the mixture of the 
black oxide of manganese and the potassium chloride so- 
lution in a beaker. Then, holding the edge of the beaker 
against the glass rod, slowly pour the contents of the beaker 
into the filter paper. The liquid should run in a small stream 
down the rod, the end of which may be held against any 
desired side of the filter paper. The liquid should not strike 
directly upon the apex of the cone. Pouring down a rod 
prevents spilling, for a liquid sometimes has a tendency 
when poured slowly from a beaker to run down the out- 
side. Pour in the liquid until its surface stands just a little 
below the top edge of the paper. If the liquid is allowed 
to go above the top edge 
of the paper there is dan- 
ger of the insoluble mat- 
ter being carried down 
in small channels that 
may exist 1 jet ween the 
filter paper and the 
ulass. AYhile the liquid 
is running thru the filter 
paper, continue adding 
the contents of the beaker 
until it has all been fil- 
tered. The liquid that 
runs thru the filter and is 
caught in the empty beak- 
er is called the filtrate. 
The insoluble substance 
left on the filter paper is 
called the precipitate. 

The filtering proceeds more rapidly if the end of the stem 
of the funnel is against the side of the beaker, because there 




Fig. 44. Filterim>;. 



52 FIRST YEAR CHEMISTRY 

is then a continuous stream of liquid pulling on the liquid 
in the filter. Warm or hot solutions filter more rapidly 
than cold solutions. In this experiment warm water was 
added to the Kjeldahl flask in order to dissolve the potas- 
sium chloride more readily; incidentally, it hastens the fil- 
tering. If any of the precipitate is left in the beaker, wash 
it out with the stream from the wash-bottle, so as to get 
all the black oxide of manganese upon the filter paper. 

Washing the precipitate. — When the liquid in the filter 
paper has all run thru, blow a stream of water from the 
wash-bottle against the sides of the filter paper and wash 
down into the apex of the filter as much of the precipitate 
as will leave the filter paper easily. Then direct the stream 
of water directly into the apex and stir up the whole mix- 
ture. By this time enough water has probably been added 
to fill the filter; if not, add enough to bring the water to 
the proper hight on the filter paper. Let this water, which 
is called the wash-water, drain thru completely. Throw 
away both the original filtrate and the wash-water. 

Drying the precipitate. — When all the wash-water has 
drained thru the filter paper, open the paper out flat so that 
it resumes its original circular form. Lay it on a piece of 
cloth or upon several thicknesses of filter paper to absorb 
most of the water. If you have plenty of time lay the 
paper away in the locker to dry of its own accord, or better, 
lay it on the radiator if the steam is turned on. If you 
are in a hurry put the opened filter paper with the precipi- 
tate on it in a clean porcelain evaporating dish, and set 
the dish on an iron gauze on the ring stand. Turn the 
flame of the Bunsen burner down to about one fourth of 
its ordinary size and have the dish at least 15 cm. above 
the top of the flame. Watch the filter paper carefully, and 
if it shows the least sign of charring, turn the flame still 
lower. When you think the moisture is all gone from the 



FIRST YEAR CHEMISTRY 53 

precipitate, weigh the paper and contents on the horn-pan 
balance. Put the paper back into the dish and heat per- 
haps a quarter of an hour longer; then weigh it as before 
and see if there has been any change in weight. If there 
has been a loss in weight it shows that all the moisture was 
not out the first time. Continue heating and weighing un- 
til it weighs twice alike; the final weight may be called 
the true weight. This weighing, heating, and weighing again 
is called heating to constant weight, and is often employed 
in quantitative chemical work. 

Getting the exact weight of the precipitate. — From the 
final weight of the filter paper and contents as just found, 
the weight of the filter paper must be subtracted in order 
to get the exact weight of the black oxide of manganese. 
This may be done very roughly by scraping off as much of 
the precipitate as possible and then getting the weight of the 
paper and subtracting this from the weight of filter paper 
and contents. A better way is to take an unused filter paper 
of the same size as the one already used, dry it to constant 
weight, and weigh it ; then use this weight in the subtraction. 
The amount of black oxide of manganese started with was 
3 grams; the amount at the end should be exactly the 
same. Did it come out so in your experiment? If there 
was any loss in weight, what do you think caused the loss? 

Purpose of the black oxide of manganese. — In the above 
experiment, the fact that the black oxide of manganese re- 
mains unchanged at the end shows that it did not enter 
into the reaction. It facilitates the chemical action during 
the breaking down of the potassium chlorate, somewhat 
"epping stones assist one in crossing a brook. This 
kind of chemical action is called catalytic action or catalysis ; 
and the substance which hastens or retards the reaction, 
but remains unchanged at the end, is called a catalytic 
agent or catalyzer. The particular advantage gained by 



54 FIRST YEAR CHEMISTRY 

the use of a catalytic agent in this experiment is that the 
oxygen is evolved more regularly and at a lower tempera- 
ture than when potassium chlorate is heated alone. We 
shall have occasion to use several catalytic agents during 
the year. In writing equations, it is always customary to 
omit the catalytic agent. 

Purity of the oxygen. — When the potassium chlorate and 
the black oxide of manganese were put into the dry Kjel- 
dahl flask and the apparatus was set up, there was some- 
thing else present in the flask, and this was carried over by 
the oxygen into the first jar that was filled. What was 
this substance? The first jar of oxygen, therefore, was not 
pure enough for the following experiments; it is pure enough 
to show the properties of oxygen. Note as many as you can 
in addition to those already found. The other jars contain 
fairly pure oxygen; save them for the next three experiments. 

Experiment 20. Action of undiluted oxygen gas on zinc. 
Have ready a Bunsen burner and a deflagrating spoon; 
also a jar of oxygen from the preceding experiment, and a 
little zinc dust. 

Put a little heap of zinc dust in the deflagrating spoon. 
Loosen the cover to the jar, but do not remove it. Have 
one hand on the cover of the jar; with the other hold the 
deflagrating spoon in the Bunsen flame until the zinc dust 
just catches fire. Then quickly but carefully lower the 
spoon into the jar and snap on the cover. Note the color 
and the intensity of the flame as compared with the burn- 
ing of zinc in air. When the combustion is complete, open 
the jar and examine the zinc oxide formed. Write the 
equation for the reaction that took place in the jar. 

Experiment 21. Action of undiluted oxygen gas on mag- 
nesium. Have ready a Bunsen burner and a deflagrat- 



FIRST YEAR CHEMISTRY 55 

ing spoon; also a jar of oxygen and some magnesium 
powder. 

Repeat Experiment 20, using powdered magnesium instead 
of zinc dust. Xote particularly the color of the oxide formed 
as compared with the color of magnesium oxide as formed 
by heating magnesium powder in air. Write the equation. 

Experiment 22. Action of undiluted oxygen gas on phos- 
phorus. Have ready a Bunsen burner and a deflagrating 
spoon; also a jar of oxygen and some phosphorus. 

Repeat Experiment 20 again, using a very small piece of 
phosphorus on the deflagrating spoon this time. Select a 
jar that has practically no water in it, use very little phos- 
phorus, and remember to do the experiment in the hood. 
Write the equation. 

Note on oxygen. — Oxygen is the most abundant of all the elements. 
It occurs in nature both free and combined with other substances. 
In the free state it occurs in the air, one fifth of which is oxygen gas. 
In the combined state it forms eight ninths of water and about one 
half of the solid crust of the earth. Many animal and plant substances 
also contain oxygen. The specific gravity of oxygen is about 1.1, re- 
ferred to air as a standard; the boiling point of liquid oxygen is — 184°C; 
its melting point has not been discovered. The principal use of oxy- 
gen gas is in the oxyhydrogen blow-pipe for getting high tempera- 
tures. A modification of oxygen, called ozone, is found in the at- 
mosphere in small quantities during thunder storms and where elec- 
tric machines are in operation; it has a peculiar odor and decomposes 
very easily into ordinary oxygen gas. 

Oxidation and kindred topics. — In Experiment 6 we learned that 
When a metal is heated its oxide is formed, and the process is called 
oxidation. We have since then made the oxides of a number of met- 
als and of several non-metallic substances. We are now in a position 
to make a general definition: Oxidation is the uniting of oxygen with 
any substance, with the formation of the oxide of that substance. It fol- 
low- naturally from this, that: An oxide is a compound substance com- 
posed of oxygen and one other simple substance. We found consider- 
able variation in the ease, rapidity, and attendant phenomena with 



56 FIRST YEAR CHEMISTRY 

which this union of oxygen with another substance took place. If 
the formation of the oxide takes place slowly at ordinary tempera- 
ture, as when iron rusts or when copper tarnishes, the process is called 
slow oxidation. When the oxidation takes place so rapidly that the 
heat generated by the reaction is sufficient to make the substance 
burst into flame, the process is called combustion. Before combustion 
can take place, it is necessary to heat the body until its kindling tem- 
perature is reached. The kindling temperature varies for different 
substances, but is always definite and constant for each one; for in- 
stance, the kindling temperature for yellow phosphorus is but little 
higher than the ordinary temperature of the air; that of zinc is some- 
what above the melting point of this metal, while that for mercury 
is entirely absent, or is so high above the temperature obtainable in 
the laboratory as to seem lacking. When any substance is heated 
until it begins to burn, the combustion will continue until the sub- 
stance is entirely oxidized, if the heat generated during the oxidation 
is sufficient to keep the substance at or above this kindling temper- 
ature; if not sufficient to do this, combustion ceases and the flame 
"goes out." By spontaneous combustion is meant combustion brought 
about when the heat caused by slow oxidation cannot escape into 
the air, but causes a rise in the temperature of the substance until 
the kindling temperature has been reached, at which point the mass 
bursts into flame. A good illustration of this is fire started in piles 
of oily and greasy cloth waste from paint shops and machine works. 

Experiment 23. The decomposition of water by means 
of electrolysis. (Lecture Experiment.) Have ready an 
electrolysis apparatus with fittings, ring stand, two clamps, 
large beaker, Bunsen burner, medium test tube, several 
wooden toothpicks, and some source of electricity, such 
as a hand dynamo or a battery; also some sulphuric acid. 

In Experiments 17 and 18, it was found that some com- 
pound substances may be decomposed by means of heat. 
Electricity instead of heat may sometimes be employed 
in a similar manner. In this experiment water will be 
broken down into the simple substances of which it is 
composed, by passing an electric current thru it. 

If the two wires from an electric battery be plunged into 



FIRST YEAR CHEMISTRY 



57 




Fig. 



45. A very simple apparatus 
for electrolyzing water. 



water in a beaker, as shown in Fig. 45, the current will 
pass from the positive pole thru the water to the negative 
pole. At the same time 



the water is decomposed 
and bubbles appear on 
both poles and rise to the 
surface. 

The electrolysis appara- 
tus shown in Fig. 46 is 
simply a modification of 
the beaker and wires men- 
tioned above; it is so arranged that the gases evolved shall 
not mix with each other. The two wires are replaced by 
two pieces of platinum foil A A; these are fastened to 
platinum wires BB, fused into pieces of glass tubing CC, 
that pass thru the rubber stoppers DD. The platinum 
wires BB are firmly fastened to copper wires EE, that 
come out from the glass tubes CC and lead to the source of 
electricity. The two graduated glass tubes F F for collect- 
ing the gases are connected by the tube G, and this in turn is 
joined to the upright tube H, which serves as a standpipe 
and carries the reservoir / for holding the excess of water. 
The upper ends of the tubes FF are supplied with glass 
stop-cocks J J. Above these are tapered glass tubes LL, 
over which rubber tubing can be slipped when necessary. 

Set up the apparatus as follows: See that the rubber 
stoppers DD are in tightly, and that the platinum foils 
are opposite the connecting tube G, and that they face 
each other. Close the stop-cocks JJ, and fill the tube H 
and reservoir / nearly to the top with water that has 
been acidified by the addition of one tenth its volume 
of sulphuric acid. Pure water is not easily decomposed 
by electricity; if a little sulphuric acid is present the wa- 
ter electrolyzes more easily and the gases are evolved 



58 



FIRST YEAR CHEMISTRY 



more rapidly; in this case the acid seems to act as a cata- 
lytic agent. Open one of the stop-cocks a little to let 

the water rise in the tube F, 
and close it when the water is 
just about to touch the stop- 
cock; be very careful not to 
let any water get into the stop- 
cock or into the narrow tube 
above it, for this is likely to cause 
considerable trouble later. As 
far as possible, no air should be 
left in the tubes FF, for the 
gases generated in the tubes F F 
are to be measured later. Fill 
the other tube F with water the 
same way, taking the same pre- 
cautions as with the first tube. 
Then connect the two wires EE 
by means of brass connectors to 
the binding posts of the hand 
dynamo shown in Fig. 47, or to 
the wires coming from the plunge 
batteries shown in Fig. 48. Then 
generate a current of electricity 
by turning the crank of the dy- 
namo or by lowering the zincs and 
carbons of the batteries into the 
solutions. If bubbles of gas are 
not generated fairly rapidly from 
the plates A A, see whether all 
the wire connections are bright. 
Collecting the gases. — Note the rapidity with which the 
two gases come off. When the tube containing the gas 
evolved the faster is about half full, disconnect the cur- 





11 , 

MM 




I 


r"WB 




■ 


H 

m 

N 

• mm 


y 






'JElSBli 





Fig. 46. A convenient elec- 
trolysis apparatus. 



FIRST YEAR CHEMISTRY 



59 




Fig. 47. Hand dynamo. 



rent and read to one tenth of a c.c. the volume of each 
gas; record these readings in the notebook; then turn on 
the current again, and let the action run until one of the 
tubes F is nearly full of gas. Disconnect the current, 
and if the batteries instead of the dynamo are used, raise 

the carbons and zincs from 

the solutions, read the vol- 
ume of each gas again, and 
record the readings; com- 
pare them with the former 
readings to see if the same 
ratio between the volumes 
of the gases still holds. 

Testing the gases. — Have 
ready a wooden toothpick 
with a glow at the end of it. Hold the glowing end ex- 
actly over the exit tube L of the tube containing the gas 

evolved in smaller quantities, 
and open the stop-cock J 
slightly, so that the inclosed 
gas may escape slowly and 
strike the glowing splinter. 
What happens to the glow? 
What one substance is thus 
proved present in water? 
Now light the Bunsen 
burner, which should stand at 
least two feet away from the electrolysis apparatus. Next 
select a medium sized dry test tube, hold it inverted over 
the exit tube L of the tube containing the gas evolved 
in larger quantities, open the stop-cock J and let the in- 
closed gas escape into the test tube till you think it has 
pushed out all the air from the test tube. Close the stop- 
cock immediately, carry the test tube still inverted over 




Fig. 48. Plunge batteries. 



60 ' FIRST YEAR CHEMISTRY 

to the Bunsen burner, and touch the mouth of the test 
tube to the flame. What happens? Watch the test tube 
closely, for sometimes the flame can be seen traveling 
slowly down the test tube. Again transfer some of this 
new gas from the tube F to the test tube, taking care not 
to displace all the air from the test tube. Now touch 
the mouth of the test tube containing the mixture of the 
gas and air to the Bunsen burner flame, and note what 
happens. Wherein does this result differ from the previ- 
ous one? This second gas obtained from water has been 
called hydrogen. What is its state, color, and solubility 
in water? Does it burn? What is the color of the flame? 
What is the test for hydrogen? Later, hydrogen will be 
made in larger quantities in more convenient ways, and 
its other properties can be studied then. 

Of what two simple substances is water composed? 
What name would you give to the chemical change tak- 
ing place when water is decomposed by means of elec- 
tricity? Where have you used this term before? Write 
the equation for the analysis of water. 

Experiment 24. The volumetric composition of water. 

(Lecture Experiment.) Have ready the electrolysis ap- 
paratus and batteries used in Experiment 23, the eudio- 
meter w r ith ring stand and a clamp to support it, a Rhum- 
korff coil, a telegraphic key, large porcelain mortar, large 
beaker, several long pieces of glass capillary tubing, and 
several rubber connectors; also ten pounds of mercury 
and a little sulphuric acid. 

Set up the apparatus as shown in Fig. 49. Clamp the 
eudiometer to its stand, and let the bottom of the U rest 
upon a block of wood or other similar support B, of such 
a hight that a large porcelain mortar can be placed under 
the outlet tube C which carries the stop cock D. To the 



FIRST YEAR CHEMISTRY 



61 



lower end of the piece of capillary tube E attach a rubber 
connector F long enough to reach around the bend of the 
eudiometer; by means of rubber connectors and a piece 
of capillary tubing G bent with two right angles, con- 




Fig. 49. Apparatus for showing the volumetric composition of water. 

nect the tube E with one of the exit tubes of the electroly- 
sis apparatus. The ends of the tube G should touch 
the ends of the glass tubes to which it is connected. Near 
the top of the closed end of the eudiometer two platinum 
wires H H are sealed into the glass in such a way that a 
slight gap exists between the two inner ends of the wires; 
the outer end of each wire forms a loop; attach each of 
these loops to the upper binding posts MM of the Rhum- 
korff coil by means of the copper wires N N. Connect 
one of the lower binding posts O of the Rhumkorff coil 
by a copper wire P to one of the binding posts of the tele- 



62 FIRST YEAR CHEMISTRY 

graphic key. Connect the other lower binding post Q 
of the Rhumkorff coil by a copper wire R to one of the 
terminals of the plunge battery. Connect the other ter- 
minal of the plunge battery by a copper wire S to the 
second binding post of the telegraphic key. All electric 
circuits are now complete, except for lowering the zincs 
and carbons into the solutions, and this is to be done at 
the last moment. Make sure that the ends of all the cop- 
per connecting wires are scraped bright and that all bind- 
ing posts are screwed down tightly so that all electric 
contacts are good. 

Generating the gases. — Now disconnect the capillary 
tube G from the electrolysis apparatus, disconnect the 
wires R and S from the battery, connect the electrolysis 
apparatus to the battery, lower the zincs and carbons 
into the solutions and generate oxygen and hydrogen 
from water slightly acidulated with sulphuric acid just 
as in the preceding experiment. Take particular pains 
not to let any liquid get above the stop-cocks JJ into 
the exit tubes LL. Generate the gases till the hydrogen 
tube is nearly full. Disconnect the electrolysis appara- 
tus from the battery and raise the carbons and zincs 
from the solutions. 

Transferring the gases. — Unclamp the eudiometer from 
the stand, close the stop cock D and, holding the tube 
carefully and firmly in the hand, fill the open arm of the 
tube with clean, dry mercury. Place the thumb over 
the open end, invert the tube, and let the mercury run 
into the closed arm of the tube, adding more mercury 
if necessary. When the closed arm is completely full of 
mercury, and no air bubbles remain clinging to the sides 
of the tube, clamp the eudiometer into position again, 
open the stop cock D and let the mercury in the open 
arm run out into the porcelain mortar. Insert the tube E 



FIRST YEAR CHEMISTRY 63 

so that the rubber tube F surely reaches around the bend 
of the eudiometer. Connect G to that tube of the elec- 
trolysis apparatus containing the oxygen. Open the stop- 
cock J and let about 15 c.c. of the oxygen pass over into 
the closed arm of the eudiometer. Close the stop-cock D, 
and fill the open arm of the eudiometer with mercury, 
till the mercury stands at the same level in both arms; 
the object of this is that the gas in the tube may be under 
the same pressure as the atmosphere of the room, other- 
wise the different volumes of gas could not be compared 
with any degree of accuracy. Read the volume of oxy- 
gen in the eudiometer, and record the reading in the note- 
book. Open the stop-cock D, and let the mercury run 
out from the open arm. Connect the tube G to that 
tube of the electrolysis apparatus containing the hydro- 
gen, and let about 15 c.c. of the hydrogen pass over into 
the eudiometer and mix with the oxygen already collect- 
ed therein. Close the stop-cock D, and fill the open arm 
of the eudiometer as before, with mercury till the mer- 
cury stands at the same level in both arms. Read the 
volume of mixed gases in the eudiometer, and record the 
reading in the notebook; record also how much of this 
last reading is hydrogen. 

Uniting the two gases. — Disconnect the tube G from the 
electrolysis apparatus, and withdraw the tube E carefully 
from the eudiometer. Add a little more mercury if nec- 
essary, so that the open arm of the eudiometer shall be 
about two thirds full. Connect the wires R and S to the 
terminals of the battery again, and lower the zincs and 
carbons into the solutions. All is now ready for the final 
touch which shall unite the two gases. This union must 
be watched for closely, as it takes place instantaneously, 
and with a slight explosion. Hold the thumb firmly 
over the open end of the eudiometer to prevent the ex- 



64 FIRST YEAR CHEMISTRY 

plosion from throwing out some of the mercury, grasp 
the telegraphic key, and, keeping the eyes on the plati- 
num wires H H and on the level of mercury in the closed 
arm of the eudiometer, complete the electric circuit by 
depressing the telegraphic key. 

Discussion of the experiment. — When the sparks jump 
across the gap between the two platinum wires, the oxy- 
gen and hydrogen unite at once. An excess of oxygen 
was passed into the eudiometer, because if just the right 
amounts of oxygen and of hydrogen were passed in, the 
force of the mercury banging up against the end of the 
eudiometer might break the tube; the excess of oxygen, 
therefore, acts as a safety cushion. Which way did the 
mercury in the closed tube move at first? What caused 
this downward movement? Release the thumb from 
the open end of the tube, and let the mercury come to rest. 
Which way does the mercury move in reaching the state 
of rest? What caused this upward movement? Finally, 
fill the open arm of the eudiometer with mercury till the 
mercury stands at the same level in both arms. For 
the third and last time, read the volume of gas in the 
eudiometer, and record the read ng in the notebook. 
Deduct the volume of this residual oxygen gas from the 
volume of the oxygen originally passed into the eudio- 
meter, and compare the volume of oxygen used with the 
volume of hydrogen passed in. What is the ratio be- 
tween these two volumes? How does this compare with 
the ratio between the volumes of oxygen and of hydrogen 
as obtained by the decomposition of water in the preceding 
experiment? The amount of liquid water formed in this 
experiment is so small that one has to look closely to find 
it in the upper part of the closed arm of the eudiometer. 
Finally raise the zincs and carbons from the battery solu- 
tions and disconnect the various pieces of apparatus. 



FIRST YEAR CHEMISTRY 65 

What happened chemically to the hydrogen when it 
combined with oxygen under the influence of the elec- 
tric spark? Write the equation for this union of hydro- 
gen with oxygen. 

Note on the Rhumkorff coil. — This instrument is composed of a soft 
iron bar or core, surrounded by a few turns of large insulated copper 
wire called a primary coil. The primary coil is put into circuit with 
the battery, the current passing thru the make and break arrange- 
ment opposite the end of the iron core. Around the primary coil, 
and entirely separate from it, is a secondary coil of very many turns 
of fine insulated copper wire, the ends of which are fastened to the 
two binding posts at the top of the coil. When the current from the 
battery is being rapidly and intermittently passed thru the primary 
coil, there is generated in the secondary coil an induced current of such 
a nature that sparks will fly across the gap between the ends of the 
wires leading from the binding posts of the secondary coil. 

Experiment 25. Preparation of hydrogen from water 
by means of magnesium. (Two students should work 
together on this experiment.) Have ready two ring 
stands, two large rings, one medium ring, one clamp, 
two Bunsen burners, two gauzes, a 500 c.c. flask, with 
cork to fit it, a narrow-necked Kjeldahl flask, pneumat- 
ic trough, and a piece of hard glass tube 20 cm. long and 
of 7 mm. internal diameter, glass delivery tube, several 
short pieces of soft glass tubing, one piece of very small 
glass tubing about 10 cm. long, several rubber connectors, 
and several old corks with holes bored thru them; also 
a piece of magnesium ribbon 30 cm. long. 

In this experiment water in the form of steam is passed 
over hot magnesium ribbon, and the hydrogen is caught 
over water in the pneumatic trough. 

Set up the apparatus as shown in Fig. 50. Fit a 500 
e.c. flask .1 with a good cork and a glass tube B bent at 
right angles; this tube should be flush with the under 



66 



FIRST YEAR CHEMISTRY 



side of the cork, and the end C should be drawn down 
to a capillary, and cut off at a such point that the di- 
ameter of the tube at C is about one half of the ordinary 
bore of the tube; over this tube B slip a fairly tightly 







Fig. 50. Apparatus for passing steam over magnesium. 

fitting cork D. Set the flask on a gauze on the large ring 
of a ring stand with a Bunsen burner underneath. Over 
the neck of the flask slip the medium ring; this is to 
keep the flask from tipping over, and still allow the neces- 
sary play when the apparatus is to be connected at the 
last moment. Fill the flask about half full of water, 
and put in half a dozen pieces of glass tubing, each be- 
tween one and two cm. long; the presence of the glass 
tubing helps in obtaining a steady stream of steam when 
the flask is heated later. E is a piece of hard glass 
tubing 20 cm. long, and of at least 7 mm. bore; into it 



FIRST YEAR CHEMISTRY 67 

put a piece of magnesium ribbon F 30 cm. long, that has 
been crumpled up loosely into a ball or coiled around a 
glass rod to form a short spiral; the ribbon should not 
be folded tightly over upon itself, for then the reaction 
does not run well. The tubes B and E should be joined 
by a rubber connector G that will slip on to B easily and 
still be fairly tight. Close to the rubber connector G, 
slip on an old, fairly tightly fitting cork H. The hard 
glass tube should be clamped near the other end to a ring 
stand which supports a Bunsen burner on a ring at such 
a bight that the magnesium in the hard glass tube may 
be heated. Inside the hard glass tube, put the piece of 
very small glass tubing /, letting it reach from the mag- 
nesium to the end of the hard glass tube; this is to keep 
the magnesium from being shot out of the tube when 
the steam is turned on. Connect the tube E with the 
pneumatic trough by means of a glass delivery tube J. 
Fill a narrow necked Kjeldahl flask L with water, and 
lay it in the trough ready for use. 

'When the apparatus is all set up as shown in the sketch, 
slip the rubber connector off from the tube B, and then 
heat the water in the flask till a small but steady stream 
of steam issues from the tube B. Put the Kjeldahl 
flask over the end of the delivery tube, then heat the tube 
E at the part containing the magnesium ribbon. From 
now on, watch the apparatus closely, for the success of 
the experiment depends upon connecting the appara- 
tus at just the right moment. Make sure that a steady, 
but not too large, stream of steam is issuing from the 
tube B. "When the magnesium begins to get red hot, 
quickly but carefully slip the tube B into the connector G, 
using the corks H and D as handles. The, burners should 
be left under the magnesium while steam is passing thru 
the tube. 'Watch the magnesium closely, for sometimes 



68 



FIRST YEAR CHEMISTRY 



the reaction takes place quickly. When the reaction has 

stopped, disconnect the apparatus at G. 

When the hydrogen has been caught in the Kjeldahl 

flask, put the finger over the mouth of the flask while 

the mouth is still 
under water, and 
take the flask out 
of the trough, 
leaving in the 
flask the water 
that was not 
pushed out by the 
hydrogen. Re- 
move the finger 
from the flask 
and touch the 
mouth of the flask 
to the Bun sen 
flame, as shown 




Fig. 51. 



Testing the gaseous product for 
inflammability. 



in Fig. 51. What happens? If the gas burns, note 
the color of the flame as the flame travels down the 
neck of the flask. What happened to the magnesium 
in this experiment? What was the color of the resi- 
due in the hard glass tube? Where did the oxygen 
come from that oxidized the magnesium? What two 
factors were there, then, in the experiment? What two 
.products were there? Write the equation for the prep- 
aration of hydrogen from steam by means of magnesium. 

This pulling of oxygen out of an oxide by means of an- 
other substance is called reduction. The substance that 
pulls the oxygen away and incidentally is itself oxidized, 
is called a reducing agent. In the above experiment, the 
steam was reduced to hydrogen, and magnesium was the 
reducing agent. Strictly speaking, a reducing agent 



FIRST YEAR CHEMISTRY 69 

must always be a material substance, generally a simple 
substance, and it must have greater attraction for the 
oxygen than the substance that is united with the oxy- 
gen in the compound to be reduced. Heat and electric- 
ity are sometimes spoken of as reducing agents, but strict- 
ly this is not correct, because we have already seen that 
when a substance has been decomposed by heat alone or 
by electricity, the change has in each case been one of 
analysis and not one of reduction. 

Note on water. — Water occurs in enormous quantities upon, in and 
around the earth. Three fourths of the surface of the globe is cov- 
ered by oceans with an average depth of two miles. There is enough 
permanent ice in the polar regions and on the tops of high mountains 
to cover all the United States with a layer of ice over one half a mile 
thick. From 50% to 90% of plants is water. The human body is 
about 70% water. Many rocks and minerals contain water chemic- 
ally combined and, therefore, not visible as water. Considerable 
water is present as moisture in the air, — enough to form a layer, if con- 
densed to liquid water, 5 inches deep over the surface of the entire 
earth. In addition to all the water mentioned above, there is an 
enormous amount in the lakes, rivers, springs, wells, and underground 
streams. The specific gravity of liquid water is 1, i.e., it is taken as 
the standard for determining the specific gravity of solids and liquids; 
water expands slightly as the temperature goes above or below 4°C, 
hence the temperature of the water must be 4° when other substances 
are compared with it; the specific gravity of ice is 0.92 referred to 
water; the specific gravity of steam is 0.62 referred to air; the melt- 
ing point of ice is 0.°C, and the boiling point of water is 100°C. Water 
dissolves very many substances, not only solids but gases and liquids 
as well. The term solubility was explained in Experiment 5, and 
the subject will be taken up again in detail when we study the theory 
of chemistry later in the year. Suffice it to say here, that all natural 
water (except rain water) contains considerable dissolved matter; 
such water may be purified, (1) by filtration to remove solid matter, 
(2) by hailing to remove dissolved gases and volatile vegetable matter, 
and (?>) by distillation to separate the water from all impurities by 
vaporizing the water and then condensing it again. When water 
freezes, the resulting ice is generally transparent and non-crystalline, 



70 FIRST YEAR CHEMISTRY 

but when water vapor in the air freezes, " snow crystals, " as they are 
called, are often formed; these are always six-sided or are built up on 
the plan of a hexagon, and vary a great deal in actual appearance. 
Figs. 52 and 53 represent a few snow crystals selected from a collection 
of several hundred photographs made by Wilson A. Bentley and pub- 
lished in the United States Monthly Weather Review. 

Experiment 26. Sulphur and its properties. Sulphur 
occurs in the laboratory in two forms, — roll sulphur and 
flowers of sulphur. Take both kinds of sulphur, and note 
as many properties of each as you can, referring to Experi- 
ment 5 to make sure you do not omit any. Compare 
sulphur with phosphorus as to properties, and find out 
which properties are more prominent in phosphorus and 
which are more prominent in sulphur. Then compare 
sulphur in the same way with the other simple substances 
already studied. 

Experiment 27. Modifications of sulphur. Have ready 
a Bunsen burner, large Hessian crucible, ring stand and 
medium ring, glass stirring rod, evaporating dish, test 
tube, test tube holder, large beaker, graduate, mortar 
and pestle, crucible cover, and magnifying glass; also 
some roll sulphur and some sulphide of carbon. 

In the preceding experiment we found two modifica- 
tions of sulphur; in this experiment we shall study three 
more. 

(A) Crush about 1 gram of roll sulphur to a fine pow- 
der in a clean mortar. Add about 5 c.c. of sulphide of 
carbon; this substance, which is one of the very few liquids 
in which sulphur will dissolve, is very volatile and in- 
flammable; therefore, keep fire away from it. Rub the 
sulphur and the sulphide of carbon together in the mor- 
tar to aid the dissolving and let the mixture stand a 
couple of minutes. Then pour the clear solution out into 




Fig. 52. Snow crystals. 




Fig. 53. Snow crystals. 



FIRST YEAR CHEMISTRY 71 

the cover of a porcelain crucible, and set it outside the 
window, so that the sulphide of carbon can evaporate 
spontaneously. When the liquid has all disappeared, ex- 
amine the form and color of the crystals deposited in the 
bottom of the cover. Use the magnifying glass in case 
the crystals are small. As far as possible, note the gen- 
eral shape of the crystals and the number of sides each 
has. Look for several crystals that are similar in shape, 
and sketch an outline of them in the notebook. In case 
the crystals did not form well, you should at least be able 
to tell whether they are long and needle-like, or short 
and chunky. The lusterless coral-like fringe of sulphur 
that sometimes forms around the edge of the dish should 
not be mistaken for sulphur crystals. If you do not re- 
member the essential characteristics of crystals, consult 
the paragraph on " How to test for properties," under 
Experiment 5. Save the crystals so as to compare them 
with other forms of sulphur obtained later. It is not 
worth while to start with flowers of sulphur in this experi- 
ment, for they do not dissolve as readily as does roll sul- 
phur. 

How to use the magnifying glass. — When examining 
the form of crystals with the aid of the magnifying glass, 
always hold the glass close to the eye and vary the dis- 
tance between the glass and the object to be examined 
till the proper focus is reached; the enlarged image of 
the object will then stand out sharply and distinctly. 

(B) Fill a large test tube from one fourth to one third 
full of powdered roll sulphur. Heat it gradually in the 
Bunsen burner flame till the sulphur melts. How does 
sulphur compare with phosphorus and with zinc in ease 
of melting? Watch the sulphur closely, as it melts, and 
note the color and thickness of the liquid. If the test 
tube gets too hot to hold with the fingers, use the test 



72 FIRST YEAR CHEMISTRY 

tube holder described in Experiment 11 on phosphorus. 
The melting point of sulphur is 114°C. Continue heat- 
ing the sulphur till it becomes thick and viscous, i.e., 
becomes sticky and of ropy or glutinous consistency; this 
takes place at about 250°C. What is its color now? 
Continue heating till the liquid actually boils; the boil- 
ing point of sulphur is 444°C. How does the condition 
of the liquid at its boiling temperature differ from that 
at its viscous temperature? As the sulphur boils does it 
vaporize, i.e., is it volatile? If it is volatile, what is the 
color of the vapor? What is the deposit on the sides 
near the top of the test tube? Scrape some of it together 
and compare it with flowers of sulphur; it is practically 
the same thing; in fact., flowers of sulphur are made on 
a large scale in just the same way. 

Now pour the hot, melted sulphur in a thin stream into 
a beaker of water. Take it out of the water and examine 
this new form of sulphur for its properties, particularly 
its hardness, color and consistency. This form of sul- 
phur is often called plastic sulphur ; sometimes it is called 
rubbery sulphur ; lay it aside in the drawer for two or 
three days and note the change that gradually comes over it. 

(C) Fill a large Hessian crucible nearly full of pow- 
dered roll sulphur. Heat it very gradually and carefully 
over the Bunsen burner; the success of this experiment 
depends upon just melting the sulphur and not letting the 
temperature run much above the melting point; if the 
sulphur becomes at all viscous good results are not likely 
to follow. Stir the melting sulphur with a glass rod to 
aid complete melting. It is a good idea, when the sul- 
phur is half melted, to take away the Bunsen burner and 
see if the rest of the sulphur will not melt of its own ac- 
cord in the heat of the liquid; if it does not, apply the 
Bunsen burner again. The sulphur, if heated carefully, 






FIRST YEAR CHEMISTRY 73 

should not catch fire. If it does catch fire, it shows that 
too much heat was applied. In that case, remove the 
burner and smother the flame of burning sulphur by set- 
ting an evaporating dish on top of the crucible. 

When the sulphur finally has been properly melted, turn 
off the gas and allow the crucible to cool. When the crystals 
have formed around the edges of the crucible and begin 
to shoot across the surface to the center, pour out what 
melted sulphur remains. The inside of the crucible should 
be coated with crystals. Examine them, particulary in 
regard to shape, brittleness and color. Compare them 
with the crystals obtained in (A). 

Note on sulphur. — Sulphur occurs free in nature in two forms, — 
crystallized and non-crystallized. The former is usually found as 
chunky crystals similar to those obtained from the sulphide of carbon 
experiment: it is called octahedral sulphur or rhombic sulphur. The 
latter, or non-crystallized, variety is simply powdered sulphur mixed 
with earth. Such native sulphur is melted to separate it from the 
earth: it is then run into molds whence it receives the name roll sul- 
phur or brimstone. \Yhen roll sulphur is heated out of contact with 
air. and the vapors of sulphur are collected as a fine powder in cool 
chambers, it is called flowers of sulphur. The flowers have no crystal- 
line form, even under the microscope; this variety is. therefore, de- 
scribed as an amorphous powder, i.e.. a shapeless powder, in order 
to distinguish it from powders that show under the microscrope that 
they are composed of minute crystals. Most free sulphur is found in 
volcanic regions. Considerable sulphur occurs in the ground combined 
with various metals; sometimes oxygen is also present. — combined, of 
course. The island of Sicily produces most of the sulphurused in the 
world: some is obtained thru manufactures; the United States deposits 
are mostly in Louisiana. The main uses of sulphur are in the manu- 
facture of gunpowder, matches, vulcanized rubber and sulphuric acid. 

Allotropy. It has just been scon that sulphur exists in crystals of 
two kinds— the prismatic or needle-like crystal and the chunky or 
rhombic crystal; it also exists in soft, plastic form, and in the two 
amorphous forms of roll sulphur and the flowers of sulphur. These 
forms are all composed of elementary sulphur. — one just as much as 
the other. One form may be changed to another, — the amorphous 



74 



FIRST YEAR CHEMISTRY 



to the crystalline and vice versa. Yet each variety has properties 
that are a little different from those of the others. Plainly the form 
of the sulphur varies; so does the color, tho not to so great an extent. 
The specific gravity is slightly different for the different varieties, tho 
it averages about 2. This existence of a substance in two or more con- 
ditions which are distinct in their relations and properties is called al- 
lotropy, and the different forms are called allotropic forms. Thus sul- 
phur has five allotropic forms. Phosphorus had two; what were 
they? Oxygen had two; what were they? Allotropy must not be 
confounded with the existence of a substance in several forms, all of 
which otherwise have similar properties. The mossy, sheet and dust 
varieties of zinc are not considered allotropic forms of zinc. 



Experiment 28. Heating sulphur in contact with air 
and in contact with oxygen. Have ready a dry Kjeldahl 
flask with cork, a delivery tube, four dry fruit jars with 
covers and washers, ring stand and clamp, Bunsen burner, 
deflagrating spoon and the hose from the bat-wing burner; 
also some potassium chlorate, black oxide of manganese, 
and roll sulphur. 

First prepare 
some oxygen. Put 
20 grams of potas- 
sium chlorate and 
5 grams of black 
oxide of manga- 
nese, intimately 
mixed, in a dry 
Kjeldahl flask. 
Do not use the 
pneumatic trough, 
but catch the gas 
in dry jars by dis- 
placement, i.e., 
pass the delivery tube, which should have a piece of 
rubber tube slipped on to the end of it, directly into 




Fig. 54. Catching oxygen by displacement of air. 



FIRST YEAR CHEMISTRY 75 

the jar. See Fig. 54. Hold the cover on as well as pos- 
sible with the hand. The oxygen gas, which is some- 
what heavier than air, will collect at the bottom of the 
jar and push the air up and out. Test, to see if the jar 
is full, by holding a glowing splinter at the crack left be- 
tween the cover and the top of the jar. When the jar is 
fvM, withdraw the rubber tube slowly, so that the oxygen 
may fill the space occupied by the tube. Snap on the 
cover at once. Set away for the following experiments 
at least three jars that are dry and full of oxygen. If 
more oxygen comes off than will fill three jars, fill as many 
more as you can, for it is always convenient to have an 
extra jar of oxygen on hand. 

Do the rest of the experiment in the hood. 

Heat a small piece of sulphur on the deflagrating spoon 
till it melts and catches fire; note the color of the flame. 
What is happening chemically to the sulphur? Do you 
see anything you might call sulphur oxide? What, then, 
is the state and color of sulphur oxide? Has sulphur 
oxide any odor? Also smell of a piece of roll sulphur 
and see if that has any odor. The residue on the spoon 
is not oxide of sulphur; generally it is only unburned 
sulphur. Make sure that the distinction between ele- 
mentary sulphur and sulphur oxide as to odor is perfectly 
clear to you. 

Set fire to a piece of roll sulphur on a clean deflagrating 
spoon, and plunge it into one of the jars of oxygen just 
made. Snap on the cover at once. Watch closely what 
goes on inside the jar. Is the color of the flame here any 
different from the flame when sulphur burns in air? What 
do you find here to be the color of sulphur oxide? Some- 
times a gray or yellow vapor is formed in the jar as the 
sulphur burns; this is not sulphur oxide; it is flowers of 
sulphur, made because some of the sulphur vaporized 



76 FIRST YEAR CHEMISTRY 

in addition to that which burned. If you get some of 
this visible vapor in your jar, let the jar stand a while till 
the vapor settles, and then get the true color of sulphur 
oxide. Get as many properties of sulphur oxide as you 
can without opening the jar. Some of the roll sulphur 
may remain unburned on the spoon, but this may be dis- 
regarded. Save the jar unopened for the next experi- 
ment. Write the equation for the burning of sulphur 
in air or oxygen. 

Experiment 29. Preparation of sulphurous acid. Have 

ready the jar of sulphur oxide from the preceding experi- 
ment, graduate, glass stirring rod, and some litmus paper. 

Take the jar of sulphur oxide from the preceding ex- 
periment. Do not open under water, but loosen the 
cover, noting whether or not there seems to be as much 
vacuum inside as there was when phosphorus was burned 
in oxygen. Remove the spoon carefully, keeping the 
cover off as short a time as possible, and add about 30 
c.c. of water. Snap on the cover again, and shake well 
in order to give the sulphur oxide a chance to dissolve, 
if it so desires. Open the jar again, and note whether 
there is any more vacuum this time. If the liquid is 
very cloudy, it might be well to filter it thru a small fun- 
nel. Taste a very small amount of the liquid in the jar. 
Try the effect of it on a picee of blue litmus paper, by tak- 
ing out a drop of the liquid on a glass rod and touching 
the drop to the litmus paper. What effect does the liquid 
have on blue litmus paper? Put a drop of distilled water 
upon a piece of blue litmus paper and see what effect 
that has. 

In the experiment for making oxygen from potassium 
chlorate, the chloride remained dissolved in the water, 
and it might have been obtained again in the solid state 



FIRST YEAR CHEMISTRY 77 

by simply evaporating the water. In this case, however, 
tho the gaseous sulphur oxide seemed to dissolve in the 
water, there was really a chemical union between the 
sulphur oxide and the water, with the production of the 
liquid, which is called sulphurous acid. What two simple 
substances are there in sulphur oxide? What two simple 
substances are there in water? If these two compounds 
unite directly, what three simple substances must there 
be in sulphurous acid? The equation for the change 
that takes place in this experiment is generally written 
as follows: 



I sulphur | j hydrogen | 

( oxygen j ( oxygen j 



/hydrogen\ 
Voxygen 
/sulphur 
[ Voxygen 



As long as both of the factors contain oxygen and there 
was a direct union between the two factors without the 
evolution of any oxygen gas, there must have been in the 
product at least two portions of oxygen. These might 
have been written together as one word, but it is custom- 
ary to write the equation as given above. The parenthe- 
ses inside the brackets in the case of the product are used 
to indicate that the compound was made from these two 
substances. The parts inside the parentheses are called 
the component parts of the product. In the above equa- 
tion, we would say, therefore, that sulphurous acid is com- 
posed of the three simple substances, hydrogen, sulphur, and 
oxygen, but that the component parts of sulphurous acid 
are sulphur oxide and water. 

Note on litmus paper. — Litmus is a vegetable compound that is or- 
dinarily blue in color, but is turned red by means of an acid and back 
to blue by means of an alkali, i.e., the opposite of an acid. Ammonia 
good example of an alkali. Litmus paper is simply filter paper 
soaked with a solution of litmus and then dried. 



78 FIRST YEAR CHEMISTRY 

Note on sulphurous acid. — On account of the ease with which sul- 
phurous acid decomposes to water and oxide of sulphur on standing 
in air, it is customary to consider sulphur oxide and sulphurous acid 
together when speaking of their uses. Their uses are in bleaching, 
fumigating and disinfecting, in preserving liquors, and in the manu- 
facture of paper and of sulphuric acid. 

Experiment 30. Heating sulphur oxide in contact with 
oxygen. (Two students should work together on this ex- 
periment.) In order to find out the general drift of the 
experiment and to ascertain on which parts of the work 
to concentrate the greatest care, read the directions 
for this experiment way thru before starting to do any 
of the work. A satisfactory result in this experiment is 
attained only by attending to details; therefore, follow 
the directions closely. 

By heating them together under the proper conditions, 
sulphur oxide and oxygen may be made to unite and form 
a second oxide of sulphur. 

Have ready two catch bottles, with good corks to fit them, 
suction pump, two fruit jars, a small battery jar, a piece 
of hard glass tubing 30 cm. long and of 7 mm. bore, tri- 
pod, gauze, Bunsen burner, bat-wing burner, two defla- 
grating spoons, about 1 m. of wash-bottle tubing, and 
several rubber connectors; also some roll sulphur, plati- 
num sponge, sulphuric acid, ice, and salt. 

First make two catch bottles, i.e., bottles that are to 
contain a liquid that will catch or absorb the impurity 
that it is desired to wash out of the air or gas used in an 
experiment. Ordinary air always has a little moisture in 
it; tho the amount is small, in such an experiment as this, 
even that small amount is detrimental to success. Select 
two wide-mouth bottles of the same size, and of perhaps 
100 c.c. capacity each. Fit these two bottles with good 
sound corks. In each cork make two holes so that glass 






FIRST YEAR CHEMISTRY 



79 



tubing of the wash-bottle size will fit it tightly. Bend 
three pieces of glass tubing, with right angle bends, so 
that when they are put in place the bottles will be con- 
nected as shown in Fig. 55. The end A of the tube B 
should dip just a little below the surface of the liquid in 
the bottle; the liquid should fill each bottle about half 
full. The end C of the tube D should be flush with the 
under side of the cork, or should protrude but little into 
the bottle; the other end E of the same tube D should 
dip just a little below the surface of the liquid in the sec- 
ond bottle. The end F of the tube G, like the end C of 
the tube D, should be flush with the under side of the cork. 
When the bottles are set up, they should be not over 5 cm. 
apart, and the horizontal parts of the tubes B, D and G 
should all be in one line. The horizontal parts of the 
tubes B and G might well be about 5 cm. in length each. 
See that all ends of the 
glass tubes are fire- 
polished. Put concen- 
trated sulphuric acid in 
each bottle so that the 
tubes dip just below the 
surface of the liquid. 
When you have finished 
using the catch bottles in 
this experiment, do not 
empty out the sulphuric 
acid, but set the catch 
bottles away just as they are, and keep them for the next 
time they arc needed. In this experiment the sulphuric acid 
removes the moisture from the air as it bubbles thru the 
acid, sulphuric acid having a great attraction for water. 

Set up the apparatus as follows: Attach to one end of 
the upright faucets a suction pump A as shown in Fig. 




Fig. 55. Catch bottles. 



80 



FIRST YEAR CHEMISTRY 



56. This suction pump is similar in principle to the filter 
pump connected with the Richards blower, described in 
Experiment 5, but it does not draw air in so rapidly, 
and is, therefore, better suited for this experiment. Bend 
a piece cf wash-bottle tubing, from 35 to 40 cm. long, 




Fig. 56. Apparatus for making the second oxide of sulphur. 

at the middle so as to form a. U; the two upright arms 
of this tube B should not be over 3 cm. apart. Bend 
each arm of this U tube about 7 or 8 cm. from the end, 
so that the ends form right angles with the body of the U. 
Immerse the tube B in a freezing mixture of ice, salt, and 
water in a battery jar. Attach the tube B to the side 
neck C of the suction pump by means of the rubber con- 
nector D. E is a piece of hard glass tubing 30 cm. long, 
and of at least 7 mm. bore; in the middle of it put some 
" platinum sponge" or " platinized asbestos," — ordinary 
asbestos fibres coated with a little powdered metallic 
platinum; the platinum sponge F should not be packed 
so tightly that the gases cannot pass thru it easily. It 
is well to insert within the hard glass tube two short pieces 
of small, soft glass tube GG — one on each side of the plati- 
num — in order to prevent the platinum from being drawn 



FIRST YEAR CHEMISTRY 81 

into the U tube by too strong suction or by uneven suc- 
tion. Eest the hard glass tube on a gauze on the tripod 
so that the platinum may be heated to a high tempera- 
ture by the Bunsen flame. Attach the tube E to the 
tube B by means of a rubber connector H; this connector 
should be as short as possible, and the ends of the two 
glass tubes inside of it should touch each other, because 
the second oxide of sulphur formed is an active substance 
and corrodes rubber. 

Drying the apparatus. — By means of the rubber tube /, 
at least 25 cm. long, connect the hard glass tube to the 
exit tube G of the catch bottles in Fig. oo: turn on the 
water a little and draw air thru the rest of the apparatus 
till all the tubes are dry inside. Considerable care §hould 
be used in this drying, otherwise the desired crystals may 
not form in the U tube, for the crystals do not form un- 
less the tubes are absolutely dry inside. Let the suction 
pump draw dry air thru the apparatus while you attend 
to the work of the next paragraph. 

Preparation of the chemicals. — While the apparatus is 
drying, prepare the factors as follows: Weigh out two 
portions of exactly 0.3 gram each of roll sulphur. Put 
one portion in a clean and dry deflagrating spoon, and 
burn the sulphur in a jar of dry oxygen as in Experiment 
28. The 0.3 gram of sulphur was not enough to use all 
of the oxygen: therefore, there was left in the jar, after 
the burning, oxide of sulphur and some free oxygen. In 
another jar of dry oxygen burn the second portion of sul- 
phur, weighing just 0.3 gram so as to be sure to have 
enough of the mixed gases. Do not remove the spoon 
from either jar. 

Passing the gases. — By this time the apparatus, thru 
which dry air has been passing, ought to be dry. Exam- 
ine the tubes closely, to < ee if it is so. Heat the plati- 



82 FIRST YEAR CHEMISTRY 

nura sponge till it gets hot. Disconnect the catch bottles, 
and put the tube / down to the bottom of one of the jars 
of mixed gases. Hold the hand over the top of the jar 
so as to prevent the air, as far as possible, from mixing 
with the contents of the jar. Let the suction pump slowly 
suck the mixed gases from the fruit jar, over the hot plati- 
num sponge, and into the cold U tube. When you think 
all the gases have been drawn from the jar, and there is 
no longer any odor of sulphur oxide in the jar, withdraw 
the rubber tube and put it into the other jar full of mixed 
gases; from this suck the gases slowly thru the apparatus 
as before. 

Study of the product. — Now look for the second oxide 
of sulphur, which should have formed a deposit of fine, 
white crystals in the bottom of the U tube; look for them 
quickly, because they melt at 14°C, which is usually but 
a little below the ordinary temperature of the air. Get 
as many of the properties of the crystals as you can with- 
out breaking the tube or emptying out any of the second 
oxide of sulphur. Describe this new compound after, as 
well as before, it has melted. 

Save the resulting liquid just as it is in the U tube for 
the next experiment. That experiment, by the way, is 
short, needing less than ten minutes for its performance. 

Write the equation for the change that took place when 
the mixed gases passed over the hot platinum sponge. 
How many portions of oxygen are there in the pro- 
duct? In this experiment the platinum sponge did not 
enter into the reaction; it acted simply as a catalytic 
agent, in that the platinum absorbed the gases and en- 
abled you to heat the condensed gases to a higher tem- 
perature than you could have heated them had you simply 
passed them thru a hot, empty tube. The platinum, 
therefore, should not be shown in the equation. 



FIRST YEAR CHEMISTRY S3 

Experiment 31. Preparation of sulphuric acid. Have 
ready the U tube containing the second oxide of sulphur, 
and the wash-bottle: also some blue litmus paper. 

Take the U tube containing the second oxide of sulphur 
from the preceding experiment, and by means of the wash- 
bottle add a few drops of water directly to the U tube. 
If some of the crystals of the second oxide are still left 
in the U tube before adding the water, listen closely when 
you add the water for the hissing sound that is often no- 
ticeable when these two substances react. The melted 
second oxide of sulphur may sometimes be changed back 
to the crystalline form, before adding the water, by im- 
mersing the U tube in the freezing mixture for a few min- 
utes. After the water has been added, try the action of 
the resulting liquid upon blue litmus paper. Add more 
water to make it more dilute and taste a little of it. The 
new substance is called sulphuric acid. Compare it in 
the two properties just mentioned with the sulphurous 
acid made in Experiment 29. Examine the concentrated 
sulphuric acid in the bottle on the shelf, and note as many 
of its properties as you can. 

Here, as in the case of sulphurous acid, there was a di- 
rect union between the water and the oxide of sulphur 
with the formation of the acid. Of course, if much water 
was added, enough of the water would be used to change 
all the sulphur oxide into the acid and then the rest of 
the water would simply dilute the acid. By evaporating 
the excess of water, naturally the pure, concentrated acid 
may be obtained. Which seems to be the stronger acid, 
sulphurous acid or sulphuric acid? 

What two simple substances are there in water? What 
two simple substances did you find in the first oxide of 
sulphur? What three simple substances did you find in 
sulphurous acid? How many component parts are there 



84 FIRST YEAR CHEMISTRY 

in sulphurous acid? What are they? How many por- 
tions of oxygen did you decide there were in the first ox- 
ide of sulphur, and how many in sulphurous acid? What 
two simple substances are there in the second oxide of 
sulphur? What are the three simple substances in sul- 
phuric acid? How many component parts are there in 
sulphuric acid? Wnat are they? How many portions 
of oxygen are there in the second oxide of sulphur? How 
many portions of oxygen are there in sulphuric acid? 
What is the difference, then, in the chemical composi- 
tion of the two oxides of sulphur? What is the difference 
in the chemical composition of sulphurous acid and sul- 
phuric acid? 

Note on sulphuric acid. — Sulphuric acid is never found free in na- 
ture. The specific gravity of sulphuric acid is 1.8; it solidifies at 
from 0°C. to 10°C, according to the amount of water present; it has 
no definite boiling point, for it decomposes when heated; its appar- 
ent boiling point is about 300°C. On account of the extreme im- 
portance of sulphuric acid, its uses are many and extensive. Enor- 
mous quantities are used in making other acids, soda, alum, glucose, 
artificial fertilizers, nitroglycerine, phosphorus, etc.; also incidentally 
in bleaching, dyeing, electroplating, and in making soap and glass. 

Historically it was known as early as the fifteenth century, being 
prepared by heating a mixture of sand and green vitriol — a compound 
containing iron, sulphur, and oxygen. The product, an oily liquid, 
was called oil of vitriol, a name still used in technical work. 

The manufacture of sulphuric acid is of considerable interest and 
importance, because this substance is the cheapest of the strong acids 
and because it is used extensively, both in the laboratory and in tech- 
nical industries. The manufacture is accomplished by two methods, 
(a) the lead chamber process and (b) the contact process. 

The lead chamber process has for centuries produced the world's 
supply of sulphuric acid. It consists essentially of passing the first 
oxide of sulphur, air, steam, and oxide of nitrogen into large cham- 
bers lined with sheet lead, — a substance not easily acted on by sulphur- 
ic acid. The sulphur oxide is generally obtained by burning sulphur, 
or some metallic ore containing sulphur, in an excess of air. The ox- 



FIRST YEAR CHEMISTRY 85 

ides of nitrogen are obtained from saltpeter. The oxides of nitrogen 
oxidize the first oxide of sulphur into the second oxide, and this with 
steam forms sulphuric acid, which collects on the walls and floors of 
the lead chambers; from here it is drained off and concentrated for 
trade. The oxides of nitrogen, which lose part of their oxygen while 
they are changing the sulphur from the first oxide to the second, are 
themselves reoxidized by the air back into the original oxides of ni- 
trogen introduced into the lead chambers; the oxides of nitrogen, 
therefore, act as carriers of oxygen, continuously taking oxygen from 
the air and giving it to the first oxide of sulphur. 

The contact process has replaced the lead chamber process in many 
quarters, tho the latter is still used to a considerable extent. The 
contact process is our laboratory method reproduced on a large scale; 
it is only recently that it has been made to work satisfactorily outside 
the laboratory. The first oxide of sulphur and air, both properly 
purified, are passed thru pipes containing the contact mixture, gener- 
ally platinum sponge. The first oxide of sulphur is oxidized to the 
second oxide and this is passed into dilute sulphuric acid or into water 
till the liquid is of the desired strength. 

Experiment 32. Preparation of hydrogen from sul- 
phuric acid, or the action of sulphuric acid on zinc. Have 
ready a 250 c.c. flask, with cork to fit it, thistle tube, 
Bunsen burner, bat-wing burner, round file, triangular 
file, ring stand with clamp and large ring, medium beaker, 

_■■ beaker, graduate, glass stirring rod, large funnel, 
filter paper, two large test tubes, small test tube, rubber 
connector, piece of hard glass tubing 20 cm. long and of 
about 7 mm. bore, and some short pieces of wash-bottle 
tubing; also some mossy zinc, sulphuric acid, and crys- 
tallized zinc sulphate. 

Making a generator. — First make a generator; this 
piece of apparatus is to be used in this and the three fol- 
lowing experiments; it will also be needed in some of the 
(experiments later in the year; s<> it i- well to keep the 
generator set up readv for use at any time. Fit a 250 c.e. 



86 



FIRST YEAR CHEMISTRY 



holes of such a size that wash-bottle tubing will fit one 
hole tightly and a thistle tube the other. Thru one hole 

of the stopper, slip a thistle tube 
so that when the stopper is in- 
serted in the flask, the lower end 
of the thistle tube will be as 
near the bottom of the flask as 
possible, without actually touch- 
ing it. Thru the other hole of 
the stopper pass a piece of glass 
tubing bent at right angles. 
The inner end of this tube should 
be flush with the lower end of 
the stopper or should protrude 
but little into the flask. The 
other arm of the right-angled 
tube should be about 10 cm. 
long. The generator, when 
finished, should look like the one 
shown in Fig. 57. 

Then set up the apparatus 
for this experiment as fol- 
lows: Take a piece of hard glass tubing 20 cm. long and 
of about 7 mm. bore. Draw it out at the middle to a 
capillary tube similar to the wash-bottle tip. The capil- 
lary part should be from 6 cm. to 8 cm. long, and should 
have a bore of 2 mm., if possible. The capillary part 
should be bent at right angles; this may be done by hold- 
ing the hard glass tube nearly horizontal and plunging it 
just for a moment into a very small blast lamp flame; if 
the capillary tube is small, it may sometimes be bent to 
advantage in the flame of a burning match. Attach the 
hard glass capillary tube A to the outlet tube B of the 
generator by means of a rubber connector. Support the 




Fig. 57. The generator ready 
for general use. 



FIRST YEAR CHEMISTRY 



87 



capillary tube by means of a ring stand and clamp if nec- 
essary. The apparatus, when finished, should be set up 
as shown in Fig. 58. 

Now generate hydro- 
gen as follows : Weigh out 
10 grams of mossy zinc; 
put it in the flask; the 
metal need be weighed 
only to within a half a 
gram. Insert the cork 
with its fittings, so that 
the end of the thistle 
tube comes near the bot- 
tom of the flask as men- 
tioned before. Put 50 
c.c. of water in a me- 
dium beaker. Measure 
out in the graduate, 
10 c.c. of concentrated 
sulphuric acid from the 
bottle on the shelf. 
Add this sulphuric acid slowly to the water in the beaker, 
and stir the mixture with a glass rod, in order to insure 
complete mixing. Sulphuric acid and water have great 
attraction for each other, and considerable heat is gen- 
erated during the mixing. It is advisable, therefore, 
alwavs to add sulphuric acid to water, and never to 
pour water into sulphuric acid, because the heat generated 
under those oircumstances is likely to cause the liquid 
to -patter around — sometimes rather vigorously; and it 
i- not comfortable 'to get sulphuric acid into the eyes; 
if this should happen, however, the best remedy is simply 
to bathe the eyes immediately with plenty of cold water. 
This mixture of acid and water is the strength ordinarily 




Fig. 58. The generator set up for 
generating hydrogen. 



88 FIRST YEAR CHEMISTRY 

used in the laboratory; it is spoken of as a strength of 
" 1:5," meaning one part of acid to five parts of water. 

Add some of this 1:5 mixture of acid and water to the 
flask containing the zinc, pouring the liquid down the 
thistle tube; at first, add just enough to cover the zinc. 
This first portion of liquid added to the flask should be 
large enough to cover the lower end of the thistle tube; 
this adding of liquid so that the end of the thistle tube dips 
under the surface of the liquid is called sealing the tube ; 
its object, of course, is to prevent the hydrogen from es- 
caping from the flask thru the thistle tube, as it would if 
this tube did not dip into the liquid. Shaking the flask 
a little helps to get the zinc down into a layer on the bot- 
tom. Bubbles of gas ought to rise very soon from the 
zinc; if they do not appear immediately, wait a few mo- 
ments for the action to start. Under no circumstances 
may the generator be heated. The action will proceed 
satisfactorily after it has once started. Add the 1:5 acid 
mixture in small portions till the whole 60 c.c. has been 
added. 

Lighting the hydrogen with the safety tube. — When you 
think the hydrogen evolved from the mixture has pushed 
all the air out of the flask, hold a small test tube inverted 
over the end of the capillary tube; when you think the 
test tube is full of gas from the flask, touch the mouth of 
the test tube to the Bunsen flame as you touched the mouth 
of the Kjeldahl flask to the flame when you prepared 
hydrogen from water by means of magnesium. The Bun- 
sen burner ought not to be near enough to the generator 
and capillary tube to run any danger of having the hy- 
drogen ignite as it issues from the capillary tube; hence, 
have it at least two feet away. If there is a distinct pop 
when the test tube is touched to the Bunsen flame, it 
shows that the hydrogen was still mixed with air; such a 



FIRST YEAR CHEMISTRY 89 

mixture explodes when brought into contact with a flame, 
and the more impure the hydrogen, — i.e., the more air 
there is mixed with it — the louder the pop. On account 
of the explosibility of impure hydrogen, the stream of 
hydrogen issuing from the capillary tube must never be 
lighted with a direct flame. Instead, fill the test tube 
with hydrogen again, and carry it still inverted to the 
Bunsen burner flame. Continue this filling and touch- 
ing to the flame until the hydrogen lights with a very faint 
pop and the blue fame of the burning hydrogen may be seen 
slowly traveling along inside the test tube as it traveled down 
the neck of the Kjeldahl flask in the experiment referred 
to a few lines back. When the hydrogen k pure enough 
to burn slowly in the test tube, carry the test tube, still 
inverted and containing the burning hydrogen, back to 
the capillary tube and touch the mouth of the test tube 
to the end of the capillary tube. The test tube should 
not be plunged far down over the capillary tube, nor held 
down there for any length of time, for that might smother 
the burning hydrogen; just touch the mouth of the test 
tube to the issuing stream of hydrogen and then take 
the test tube away. The hydrogen should now burn 
quietly at the end of the capillary tube. This method 
of lighting hydrogen gas is called lighting it with the safety 
tube, and this method must always be used in order to avoid 
an explosion which might throw broken glass around. 

Study of the hydrogen flame.— Note the color of the 
hydrogen flame to see if it agrees with the color of the 
hydrogen flame as seen in Experiment 25. When the 
hydrogen burns, what is happening to it chemically? 
What is the chemical name of the product formed when 
hydrogen oxidizes? Tan you see the product? Hold a 
fold, dry beaker over the hydrogen flame for a moment, 
and note the deposit formed on the inside of the beaker. 



90 FIRST YEAR CHEMISTRY 

Rub the finger over the deposit in order to collect some 
of it. What is it? What is the common name for oxide 
of hydrogen? Write the equation for what happens 
when hydrogen burns. 

Note on the color of the hydrogen flame. — Occasionally, when hy- 
drogen burns at the end of a glass tube the flame is tinged more or 
less with yellow. Absolutely pure hydrogen burns with a pale blue 
flame, but this result is usually hard to obtain with ordinary appara- 
tus. The yellow flame, referred to above, is due to the presence of a 
small amount of soda in the glass. This soda is introduced as one 
of the ingredients at the time the glass is made, and the heat of the 
burning hydrogen vaporizes a very small amount of this soda. The 
yellow color is always imparted to a flame if soda is present, and a 
very small amount of soda is sufficient to give considerable yellow 
to the flame. If the laboratory is supplied with a glass tube having 
a small tube of platinum foil fused into the end of it, this platinum- 
tipped tube may well be used to show the color of the flame when pure 
hydrogen burns. Platinum has no soda in it as does the glass, and 
no soda can get into the hydrogen itself if the zinc and sulphuric acid 
are of a reasonable degree of purity. 

Treatment of the residue in the flask. — When you have 
played with the hydrogen flame as long as you wish, blow 
out the flame, remove the stopper and the fittings from 
the flask, and empty the contents of the flask into a beaker. 
Be sure to get out all the zinc as well as the liquid, for 
the amounts of zinc and sulphuric acid taken were such 
that when all the zinc has disappeared there will be just 
a little free sulphuric acid left in the solution — enough 
to give good crystals, but not enough to do any harm in 
the subsequent treatment of the solution. Let the beaker 
stand till the zinc has been entirely eaten up by the acid; 
this ought not to take over twenty minutes or half an 
hour. If you are in a hurry, set the beaker on the gauze 
on the tripod and heat gently, watching it carefully, lest 
the mixture froth over the top of the beaker; it is safe, 



FIRST YEAR CHEMISTRY 91 

of course, to heat the hydrogen-producing mixture in an 
open beaker, tho it was not allowable to heat it when it 
was confined in the generator. When all action has 
stopped, if there is any sediment or black particles in the 
liquid, filter thru a large funnel, using the precautions 
for filtering given in the experiment where oxygen was 
made from potassium chlorate and black oxide of man- 
ganese. Catch the filtrate in a clean beaker. 

Crystallization. — Set the clear filtrate away in the bot- 
tom of the locker, and let it stand till a deposit of crystals 
forms on the bottom of the beaker. If directions have 
been followed carefully, these crystals ought to appear in 
about three hours, or possibly a little less; this will depend, 
of course, upon the strength of the solution, the more di- 
lute the solution the longer the wait till crystals appear. 
If no crystals appear after three hours, evaporate the 
liquid to about half its volume. A better method is to 
evaporate the liquid till a few c.c. of the boiling liquid 
taken out in a test tube and cooled under the water tap 
deposit small crystals. Then set it away. This is called 
evaporating to crystallization. Ordinarily, the slower the 
crystals are in forming the larger will they be. 

When a good deposit of crystals is formed, filter, and 
spread the crystals out on a dry filter paper for a few min- 
lites, in order to remove all liquid from the crystals. Get 
as many properties as you can of the crystals, particu- 
larly their shape, color, luster, transparency, taste, and 
solubility in water. If the crystals are small, examine 
them under the magnifying glass. 

Composition of the crystals. — Sulphuric acid has what 
three simple substances in it? When you made sulphuric 
acid by adding water to the second oxide of sulphur you 
found that the acid contained three different portions of 
oxygen; in writing the equation for that change it was 



92 FIRST YEAR CHEMISTRY 

proper to show these three separate portions of oxygen 
in the product. It is neither necessary nor advisable, 
however, to show these three separate portions of oxygen 
every time sulphuric acid is used. A large amount of 
work done on sulphuric acid by other experimentors has 
shown that the oxygen in sulphuric acid is all very tightly 
united to the sulphur, and that in case of the decomposi- 
tion of the acid by means of another substance, it almost 
always happens that the hydrogen is liberated and that 
all the oxygen stays with the sulphur and that this com- 
bined group of sulphur and oxygen acts as a unit and unites 
with any metal that happens to be present. If the sub- 
stance that decomposes the acid is a metal, the hydrogen 
is generally replaced by that metal, and the substance 
formed by evaporation of the solution is composed of the 
metal, sulphur, and all three portions of oxygen. The 
equation for the action of sulphuric acid on zinc is, there- 
fore, written as follows: 

■ r 1 ( hydrogen ") , ^ I zinc ^ 

-j zinc [■ -\- < sulphur [- = < hydrogen j- -f- -j sulphur >- 
*- (_ oxygen J ( oxygen J 

In the above equation it is generally understood that 
the oxygen that is written only once in the sulphuric acid 
contains the three portions of oxygen mentioned above; 
in like manner, the oxgyen written once in the product 
containing zinc, sulphur, and oxygen, contains the same 
three portions of oxygen. The compound containing zinc, 
sulphur, and oxygen is called zinc sulphate or sulphate of 
zinc. What are the three simple substances in zinc sul- 
phate ? 

When the crystals are perfectly dry put a few of them 
in a clean and dry test tube, and heat slightly, noting 
what happens. Does a deposit form on the sides of the 



FIRST YEAR CHEMISTRY 93 

test tube'? What does it look like? Examine some of 
the zinc sulphate from the bottle on the shelf and see if 
it differs any from the zinc sulphate you made. Heat a 
few of the crystals of it in a dry test tube and see if you 
get the same deposit of moisture on the sides of the test 
tube, that you got when you heated the zinc sulphate you 
made yourself. This moisture is contained in the crystals 
of sulphate of zinc; it is thought to be chemically united 
with the zinc, sulphur, and oxygen, but the union is not 
verv strong, as is shown by the fact that heat drives the 
water out readily. 

The white residue left in the test tube contained all 
the zinc, the sulphur, and the oxygen considered in the 
above equation. The water that was driven out of the 
crystals is called water of crystallization. The formula 
for zinc sulphate in the equation did not show the pres- 
ence of this water of crystallization, because that equation 
showed the formation of hydrogen, and under those cir- 
cumstances the zinc sulphate Avas in solution and was not 
present in the crystallized state. The exact form or 
state in which zinc sulphate exists in the solution is not 
known definitely; it has been agreed, however, that when- 
ever an equation represents a substance like zinc sulphate 
in solution, this water of crystallization that w T e found in 
the crystals shall never be represented. 

Suppose, however, that we want to express on paper 
the fact that water of crystallization is driven out from 
the crystals by heat; in that case we would write the 
equation thus: 

r/zinc \ 1 

sulphur C zinc ) , , , . 

WW \ = -sulphur + s:^ en 

| /hydrogen \ (oxygen ) (oxygen j 



/hydro non v 
v oxygen / 



I ^oxygen 



94 FIRST YEAR CHEMISTRY 

Here we see two compounds, each containing zinc, sulphur, 
and oxygen; the first one contains water of crystalliza- 
tion, while the second one does not. To distinguish be- 
tween the two, the first one is called crystallized zinc sul- 
phate, and the second one is called amorphous zinc sul- 
phate or better, anhydrous zinc sulphate, i.e., waterless 
zinc sulphate. The white residue left in the test tube 
was the anhydrous sulphate of zinc; it had, of course, 
lost its water of crystallization, and at the same time its 
crystalline form. From this it would seem that crystal- 
line form indicates the presence of water of crystalliza- 
tion; it does in some cases; crystals containing water of 
crystallization are generally made more readily than crys- 
tals which do not contain water of crystallization. A 
number of waterless crystals do occur, however; elemen- 
tary sulphur in some of its allotropic forms is a notable 
example. 

Effect of air on zinc sulphate. — Select several good, dry, 
clear crystals of zinc sulphate; allow them to lie exposed 
to the air for several hours, or better, for several days; 
it is sufficient to leave them inside the locker. Examine 
them from time to time to see what change comes over 
them. When they have entirely crumbled to a white 
powder, the same reaction has taken place that took place 
when you heated the crystals in the test tube, but it has 
taken place more slowly. In other words, the crystals 
have lost their water of crystallization and have been 
changed to anhydrous zinc sulphate. The loss of water 
of crystallization during exposure to air is called efflores- 
cence. Do not confound this with effervescence, which is 
the rapid bubbling of a gas out of a liquid on exposure to air. 
Whenever any crystals of a new compound are made, 
you should always determine whether or not the crystals 
contain water of crystallization; and if they do, see if 



FIRST YEAR CHEMISTRY 95 

the crystals are efflorescent on long or short standing 
in air. 

What are the three simple substances in sulphuric acid? 
What are the three simple substances in anhydrous zinc 
sulphate? What are the four simple substances in crys- 
tallized zinc sulphate? What are the two component parts 
in crystallized zinc sulphate? Are these two component 
parts tightly united or not? How does the intimacy of the 
union between the component parts of crystallized zinc 
sulphate compare with the intimacy of the union between 
the three simple substances in anhydrous zinc sulphate? 

Note on zinc sulphate. — Since sulphuric acid, or oil of vitriol, as it 
is often called, is such a common substance, it follows that the sul- 
phates are frequently met with. Zinc sulphate is no exception to this 
rule; it has been called white vitriol, partly on account of its color 
and partly on account of its connection with oil of vitriol. 

Note on sulphates in general. — Any substance made by the replace- 
ment of hydrogen in sulphuric acid by a metal is called a sulphate of 
that metal. The sulphates are common substances, and many of 
them have fairly definite crystalline form; it will be of interest, there- 
fore, to make a number of sulphates before leaving the subject of 
sulphur and its compounds. 

Experiment 33. Action of sulphuric acid on iron. Have 

ready a generator, the hard glass capillary tube used in 
the preceding experiment, Bunsen burner, ring stand with 
clamp and large ring, gauze, medium beaker, large beaker, 
graduate, glass stirring rod, large funnel, filter paper, 
two large test tubes, and a small test tube; also some 
iron filings, sulphuric acid, and crystallized iron sulphate. 
In this case, generate hydrogen as follows: Put 10 grams 
of clean iron filings in the generator flask; the filings 
should be weighed to within one half a gram. Insert the 
cork and fittings properly. Put 50 c.c. of water in a medi- 
um beaker; measure out into a graduate 10 c.c. of concen- 



96 FIRST YEAR CHEMISTRY 

trated sulphuric acid and add it slowly to the water with 
enough stirring to insure mixing of the liquids. Add a 
little of this mixture to the flask containing the iron, 
pouring the liquid down the thistle tube in just the same 
way as in the experiment on zinc and sulphuric acid. Be 
sure that the first portion of the 1 : 5 acid mixture is large 
enough to seal the thistle tube. The action in this case 
between the metal and the acid is often not quite so vig- 
orous as in the case of zinc and sulphuric acid; if the ac- 
tion does not start readily, add some more of the diluted 
acid from the beaker and let the generator stand till the 
action begins; you ought not to have to wait more than 
10 or 15 minutes at the longest. Add the 1:5 acid mix- 
ture in small portions till the whole 60 c.c. has been added 
and shake after each addition. 

Study of the hydrogen flame.— When the gas is being 
evolved in sufficient quantities, test the gas issuing from 
the capillary tube, using the test tube method explained 
in the preceding experiment. Finally light the hydrogen 
with the safety tube. Note the color of the hydrogen flame, 
comparing it with the color of burning hydrogen that you 
have seen before. When the hydrogen burns, what is 
happening to it chemically? What is the product of the 
combustion of hydrogen? As before, hold a cold, dry 
beaker over the hydrogen flame and note the deposit 
formed on the inside of the beaker. Write the equation 
for the burning of hydrogen. 

Treatment of the residue in the flask. — When you have 
proved that hydrogen is produced from sulphuric acid by 
means of iron, empty the contents of the flask into a large 
beaker, being careful to transfer all the iron to the beaker. 
Let the reaction between the two factors run to an end; 
as it may take from three to four hours for the acid to 
eat up the iron completely, either set the beaker away 



FIRST YEAR CHEMISTRY 97 

in the locker over night, or heat the beaker gently on an 
iron gauze over a Bunsen burner; the liquid should not 
be actually boiled, for this would evaporate some of the 
water, with the result that the concentrated acid solution 
might not act as readily upon the remaining particles of 
iron. When the reaction has practically stopped, filter 
the liquid from any black particles floating in it; these 
particles are most likely carbon, and other impuri- 
ties always found in commercial iron. The filtrate should 
be perfectly clear, but green in color, and with a some- 
what acid reaction. Evaporate the liquid to about one 
half its volume, and set it away to crystallize, or better 
evaporate to crystallization, as described in the preced- 
ing experiment. If any white sediment appears in the 
liquid during the boiling, stop heating immediately and 
set the liquid away to crystallize. If the solution is con- 
centrated and not too acid, the crystals may appear as 
rapidly as did the zinc sulphate crystals; but if the solu- 
tion is too dilute or too strongly acid, the crystals wanted 
here may not appear for a long time, — perhaps not for 
several days. 

When a good deposit of crystals has formed on the bot- 
tom of the beaker, filter from the remaining liquid, and 
spread the crystals out on filter paper to dry for a few 
minutes in order to remove all liquid from them. Get as 
many properties as you can of the crystals, particularly 
their shape, color, luster when freshly made, transpar- 
ency, taste, and solubility in water. In what two respects 
are these crystals markedly different from those of crys- 
tallized sulphate of zinc? 

Composition of the crystals. — When sulphuric acid acted 
on zinc, what compound of zinc did you finally find that the 
resulting liquid contained? When sulphuric acid acted 
on iron in this experiment with the evolution of hydro- 



98 FIRST YEAR CHEMISTRY 

gen, what compound of iron did the resulting green solu- 
tion probably contain? Write the equation for the re- 
action between sulphuric acid and iron, referring, in case 
of doubt, to the corresponding equation in the preceding 
experiment. Here, as in the case of zinc, it is best to 
write the formula for the resulting iron sulphate as con- 
taining simply iron, sulphur, and one portion of oxygen. 

Select two or three good, dry, clear crystals of sulphate 
of iron from those you made, and allow them to lie ex- 
posed to the air for several days to see if they effloresce. 
Select again two or three equally good crystals of crys- 
tallized iron sulphate from those you made, put them in 
a dry test tube, and heat gently to find out if they con- 
tain any water of crystallization. If you find a deposit 
of moisture on the sides of the test tube, examine the resi- 
due to discover the color of anhydrous iron sulphate. If 
the residue is brown, it shows you used too much heat. 

Examine some of the sulphate of iron from the bottle 
on the shelf for all the properties that you found in your 
own sulphate of iron; be sure that you include the tests 
for water of crystallization and for efflorescence, also the 
color of crystallized and of anhydrous sulphate of iron. 
The sulphate of iron found in the bottle on the shelf will 
probably not have any definite crystalline shape, because 
it was made on a large scale, and the pieces we get in trade 
are fragments broken from large, thick deposits. Your 
own crystals ought to show fairly definite crystalline 
form. 

What are the three simple substances in anhydrous 
iron sulphate? What are the four simple substances in 
crystallized iron sulphate? What are the two component 
parts of crystallized iron sulphate? Are these two com- 
ponent parts tightly united or not? Write the equation 
for driving water out of crystallized iron sulphate. 



FIRST YEAR CHEMISTRY 99 

Substitution. — This experiment, in which a simple sub- 
stance (iron) replaced a simple substance (hydrogen) in 
a compound, and the preceding experiment, in which a 
simple substance (zinc) replaced a simple substance (hy- 
drogen) in a compound, are samples of the kind of chem- 
ical change called substitution. Substitution is the replace- 
ment of a simple substance in a compound by another simple 
substance capable of reacting with the compound; the simple 
substance replaced is usually set free in its elementary 
state. 

Note on iron sulphate. — What is the common name for crystallized 
zinc sulphate? "What is the common name for sulphuric acid from 
which this common name for zinc sulphate is derived? Crystallized 
iron sulphate is also a vitriol, and on account of its color is called 
green vitriol ; this name is very often used for the crude compound. 
The name copperas is occasionally used to denote the same substance. 

Experiment 34. The action of sulphuric acid on mag- 
nesium. Have ready the generator and capillary tube 
used before, Bunsen burner, ring stand and large ring, 
gauze, medium beaker, large beaker, graduate, glass stir- 
ring rod, large funnel, filter paper, two large test tubes, 
and a small test tube; also some magnesium, sulphuric 
acid, and crystallized magnesium sulphate. 

This experiment is entirely similar to the action of sul- 
phuric acid on zinc and on iron, but it is easier to per- 
form and it takes less time. Either form of magnesium 
may be used. 

Again generate hydrogen as follows: Put a definite 
weight of magnesium in the generator flask. If you use 
the powder, weigh out exactly 1 gram; for use with the 
powder make up a solution in a beaker of 100 c.c. of water 
and 5 c.c. of acid. If you use the ribbon, take a piece of 
magnesium ribbon about 60 cm. long, and for this make 



100 FIRST YEAR CHEMISTRY 

up a solution of 20 c.c. of water and 1 c.c. of sulphuric 
acid. More satisfactory crystals are generally obtained 
from magnesium ribbon than from the powder. Test the 
gas evolved from sulphuric acid by means of magnesium 
in the same way that you tested for hydrogen in the two 
preceding experiments, i.e., test with the safety tube. When 
the reaction has run to an end, empty the contents of the 
generator into a beaker, boil it down to about one half its 
volume, or better evaporate to crystallization, and set it 
away to crystallize. The crystals ought to appear in the 
course of a few hours, but the exact time naturally de- 
pends upon the strength of the solution. 

When a good deposit of crystals has formed, filter and 
dry the crystals on filter paper. Get as many properties 
as you can of the crystals, comparing them with the sul- 
phates of zinc and of iron. 

Composition of the crystals. — What name would you 
give to the crystals obtained in this case when sulphuric 
acid acts on magnesium? Write the equation for the ac- 
tion of sulphuric acid on magnesium. Heat some of the 
dry, crystallized sulphate of magnesium that you made, 
and see whether it contains water of crystallization. Also 
expose some of the crystals to air for several days to find 
out whether they effloresce. Write the equation for sep- 
arating water of crystallization from crystallized magne- 
sium sulphate by means of heat. Compare the sulphate 
in the bottle on the shelf with that which you made; try 
particularly the test for water of crystallization and note 
the color of the residue. What are the three simple sub- 
stances in anhydrous magnesium sulphate? What are the 
four simple substances in crystallized magnesium sul- 
phate? What are the component parts of crystallized 
magnesium sulphate? Are these two component parts 
tightly united or not? 



FIRST YEAR CHEMISTRY 101 

Note on magnesium sulphate. — This sulphate, like the other two 
just studied, is a very common substance. Crystallized magnesium 
sulphate is often called Epsom salt, because this substance was origi- 
nally found in the mineral springs at Epsom, England. 

Definition of the term "salt". — The common name mentioned in the 
preceding paragraph introduces a new term, — salt. In ordinary lan- 
guage "salt" indicates the white powder used in seasoning food. Re- 
garded in the light of its chemical composition, table salt is to the 
chemist similar to a large number of other compounds, all of which 
may be made by replacing the hydrogen of an acid by means of a 
metal. In this light all such compounds may be considered salts. 
A salt, therefore, is a compound formed by the replacement of the hy- 
drogen of an acid by a metal. The sulphates of zinc, of iron, and of 
magnesium are all considered salts of sulphuric acid. 

Experiment 35. Action of sulphuric acid on copper. Have 

ready the generator used before, Bunsen burner, ring stand 
with clamp and large ring, gauze, medium beaker, large 
beaker, graduate, glass stirring rod, large funnel, filter paper, 
several large test tubes, long glass delivery tube, and pneu- 
matic trough; also some copper turnings, sulphuric acid, 
and crystallized copper sulphate. 

It was stated some time ago, that when sulphuric acid 
is put on a metal, hydrogen is usually evolved; the three 
preceding experiments proved this true. There are, how- 
ever, a very few exceptions, copper being the best known 
of these. Let us study the action of this metal with sul- 
phuric acid. 

Generation of gas. — Put 10 grams of copper turnings in 
the generator flask, insert the stopper with its fittings, but 
in place of the hard glass capillary tube use a glass deliv- 
ery tube reaching to the pneumatic trough. Be sure that 
the end of the thistle tube is near the bottom of the flask, 
then add 25 c.c. of concentrated sulphuric acid to the flask; 
this ought to be enough to seal the thistle tube; if it is not, 
add a little more concentrated sulphuric acid, — just enough 



102 FIRST YEAR CHEMISTRY 

to seal the tube. Set the flask on the gauze and bring the 
mixture to a boil. Keep it boiling gently for about 10 or 
15 minutes, and from time to time catch the gas evolved 
in test tubes over the pneumatic trough. If it does not 
give the test for hydrogen, find out by its odor, or other- 
wise, what gas it is. Where have you made this gas be- 
fore? What simple substances are there in this gas? 

Treatment of the residue in the flask. — When the mix- 
ture has been boiled for 10 or 15 minutes, disconnect the 
apparatus and stop heating. The residue in the flask should 
now be a pasty mass, but if the mixture was boiled too 
vigorously, it may have gone just to dryness; in either 
case, let the flask cool and then add very cautiously about 
25 c.c. of water. Shake well and let stand for several min- 
utes to insure complete solution of the product; warm a 
little if you think it necessary; then empty the contents 
of the flask into a beaker. In this case do not allow any 
pieces of copper to remain in the liquid. To separate the 
liquid from this sediment, pour the liquid slowly into an- 
other beaker, leaving the sediment behind in the first beak- 
er. This separating a liquid from a precipitate or sediment 
by simply pouring it off is called decantation ; the liquid 
that is poured off is called the decantate ; this method is al- 
ways used instead of filtering in the case of a precipitate 
that settles well, or in the case of a solution that — like this 
one — can not be filtered ; in this particular instance the solu- 
tion is strongly acid, and it would, therefore, decompose the 
filter paper. Set the clear decantate away to crystallize; this 
may take several days on account of the strong acid present. 

When a deposit of crystals has finally formed, separate 
them from the liquid, dry them, and get their properties 
as in the case of the three sulphates just studied. Com- 
pare your crystals with some crystallized sulphate of cop- 
per from the bottle on the shelf. 



FIRST YEAR CHEMISTRY 103 

Composition of the crystals. — The reaction in this experi- 
ment was rather complicated; still, it is possible to write 
an equation. Note that the first oxide of sulphur, and not 
hydrogen, was evolved, and that sulphate of copper was 
left in the solution. Therefore, two separate portions of 
sulphuric acid are needed, the copper taking the sulphur 
and oxygen from one portion of the acid, forming copper 
sulphate, while the liberated hydrogen takes some of the 
oxygen from the other portion of sulphuric acid, forming 
water and setting free the oxide of sulphur. In writing 
the equation, insert sulphuric acid twice; write the first 
one as you have written sulphuric acid in the other preced- 
ing equations; write the second one as you wrote sulphuric 
acid in Experiment 31, so as to show all the separate por- 
tions of oxygen. Heat some dry crystals of crystallized 
copper sulphate in a dry test tube, to see if they contain 
any water of crystallization. What is the color of anhy- 
drous copper sulphate? Allow a good, clear crystal to lie 
exposed to the air for several days to see if it effloresces. 

What three simple substances are there in anhydrous 
copper sulphate? What four simple substances are there 
in crystallized copper sulphate? What are the two com- 
ponent parts of crystallized copper sulphate? Are these 
two component parts tightly united or not? Write the 
equation for separating the component parts by means of 
heat. 

Recrystallization. — Large, good crystals of copper sul- 
phate may be obtained by dissolving some crystallized 
copper sulphate in water and allowing the salt to crystal- 
lize -lowly. Weigh out 50 grams of copper sulphate from 
the bottle on the shelf, dissolve it in 100 c.c. of water in a 
beaker by means of heating. Set the liquid away to crystal- 
lize. When the crystals have formed, examine them for shape, 
size, and color, comparing them with those already made. 



104 FIRST YEAR CHEMISTRY 

Note on copper sulphate. — We have found two salts called vitriols, 
and they were named after their colors. What were they? Crystal- 
lized copper sulphate belongs to the same class, being called blue 
vitriol ; it is occasionally called blue stone. It is used extensively 
in electric batteries, in copper plating, and in making other com- 
pounds of copper. There are no other vitriols that are named after 
their colors. 

Experiment 36. The action of sulphuric acid on zinc ox- 
ide. Have ready the generator and capillary tube used 
before, Bunsen burner, ring stand with clamp and large 
ring, gauze, two beakers, graduate, glass stirring rod, large 
funnel, filter paper, and two large test tubes; also some 
zinc oxide and sulphuric acid. 

Put 10 grams of oxide of zinc from the bottle on the shelf 
in the generator flask. In a beaker make a solution of 30 
c.c. of water and 6 c.c. of concentrated sulphuric acid. Add 
this mixture to the flask and wait a little while to see if 
any action begins. If any gas is evolved in large quanti- 
ties, test it for hydrogen. If no gas is evolved in five or 
ten minutes, shake the flask well and heat it gently for some 
time. Keep the mixture near the boling point, but do not 
actually boil it. If any gas is evolved now, test it for hy- 
drogen. If you can find no hydrogen, see if you can find 
out what became of the hydrogen. In order to determine 
this, filter the mixture in the flask, evaporate the clear fil- 
trate to about half its volume, and set it away to crystallize. 
When the crystals have formed, see if you can tell by their 
shape, <jolor, and any other properties where you have met 
this substance before. Heat one of the dry crystals in a dry 
test tube to see if any water of crystallization is present. 
How many factors did you start with? Was either factor 
a simple substance? Of what simple substances was each 
factor composed? What were the three simple substances 
in that product which later crystallized out from the solu- 



FIRST YEAR CHEMISTRY 105 

tion? What simple substances are now left unaccounted 
for in the two factors? There were only two products in 
this reaction, and neither was a simple substance. What, 
then, must the other product have been? Write the equa- 
tion for the change that takes place when zinc oxide is treat- 
ed with sulphuric acid. 

Note on the action of sulphuric acid on oxides. — Sulphuric acid re- 
acts with the oxides of iron, magnesium, and copper in the same way 
that it acted on oxide of zinc, but we shall not take the time to 
perform these experiments. In each case the sulphate of the 
metal would be formed and water would be the second product. 
Generally speaking, sulphuric acid reacts with all oxides in a similar 
manner. 



Experiment 37. Reaction between copper and sulphur. 

Have read}- a test tube, Bunsen burner, ring stand with 
clamp and small ring, crucible tongs, and mortar and pes- 
tle; also some roll sulphur and some copper sheet and wire. 

Do both parts of this experiment in the hood. 

In Experiments 6, 8, 10, and 14 we studied the action of 
oxygen upon the metals copper, zinc, magnesium, and iron. 
We shall now study the action of hot sulphur upon several 
metals, beginning with copper. 

Fill the test tube about one fourth full of roll sulphur; 
clamp it, about one third down from the mouth of the tube, 
at such a hight that the lower end of the tube may be heated 
by the Bunsen burner; the test tube should be held at an 
angle of about 45 degrees. 

Heat the sulphur till it boils vigorously. When the va- 
pors of sulphur rise to the mouth of the test tube, touch a 
lighted match to them. The vapors should catch fire and 
burn with a blue flame at the mouth of the tube. Thru- 
out the whole experiment this flame should extend not more 
than three or four cm. above the mouth of the test tube. 



106 FIRST YEAR CHEMISTRY 

Roll a piece of coarse, copper wire around a pencil to make 
a little coil of it. Pull this coil out to a length of a couple 
of centimeters, and then hold it with the tongs in the flame 
of the burning sulphur for a few minutes; hold it in such a 
way that part of the copper is directly in the flame of the 
burning sulphur and part of it is down in the test tube be- 
hind the sulphur flame and in contact with the hot vapors 
of sulphur. What change do you notice in the appear- 
ance of the copper? Let the wire cool; then scrape off the 
black coating on the outside of the wire and get its proper- 
ties. Did any coating form on that part of the wire that 
was in contact with the sulphur vapors only? Could there 
have been any oxygen in that part of the test tube? Is 
there any oxygen, then, in the coating? 

There has been a direct union between the copper and 
the sulphur in this case, and the oxygen did not enter into 
the reaction. The product is called copper sulphide or sul- 
phide of copper; and it contains two simple substances, 
copper and sulphur. The equation is: 

copp e r } + { sulphur } = { "g*^ } 

Note the similarity between oxygen and sulphur in their 
behavior with hot copper. 

Get a larger amount of copper sulphide as follows: Take 
a piece of sheet copper, hammer it out on the anvil till it 
forms a shallow spoon, and support this on the small ring 
of the ring stand, over the naked Bunsen burner flame. 
Put a small piece of roll sulphur in the hollowed copper 
and heat the copper sheet till the sulphur has stopped burn- 
ing. Let it cool and then detach the scale of copper sul- 
phide. Compare this scale with the copper sulphide al- 
ready made. Reduce it to a powder in a mortar. Save the 
product for Experiment 44. 



FIRST YEAR CHEMISTRY 107 

Experiment 38. Reaction between mercury and sulphur. 
Have ready a mortar and pestle; also some mercury and 
some flowers of sulphur. 

Put a small globule of mercury, not over 2 cm. in diameter, 
in a clean and dry mortar. Add to it about twice its vol- 
ume of flowers of sulphur. Grind the mixture for several 
minutes, or till you think the union is complete. This 
process is called trituration and is employed occasionally 
in chemical work. To triturate means to pulverize and 
mix thorohj by grinding. If any unchanged mercury re- 
mains, separate it from the resulting powder. Get the prop- 
erties of mercury sulphide. Save the product for Experi- 
ment 44. Write the equation for the formation of mer- 
cury sulphide. Which forms more readily, i.e., at lower 
temperature, copper sulphide or mercury sulphide? 

Experiment 39. Reaction between zinc and sulphur. Have 

ready a Bunsen burner, asbestos sheet, iron rod, and spatu- 
la; also some zinc dust and flowers of sulphur. 

Do this experiment in the hood. 

Take about 10 c.c. of zinc dust and an equal volume of 
flowers of sulphur. Mix the two intimately on the asbestos 
sheet by means of the spatula till the mixture is of an even 
color and contains no lumps of either factor. Scrape the 
mixture into a little heap on the asbestos sheet and place 
it in the hood. Heat an iron rod in the Bunsen flame till 
it is red hot and then touch it to the mixture on the asbestos 
sheet. Note what happens. Note also the color of the 
smoke. Note also the color of the residue on the asbestos 
sheet while hot and watch the change in color as the residue 
cools. Examine the residue and get as many properties 
as you ran of zinc sulphide, Write the equation for the 
uniting of zinc with sulphur. A little of the zinc or of the 
sulphur, of or both, may have oxidized during the forma- 



108 FIRST YEAR CHEMISTRY 

tion of the sulphide, but this may be neglected and need 
not be shown in the equation. Save the residue for Experi- 
ment 44. How does the union between zinc and sulphur 
differ from that between mercury and sulphur or that be- 
tween copper and sulphur? In which case was the action 
most vigorous? In which case did the action proceed of 
itself when it was once started? In which case did the 
uniting begin with practically no outside aid? 

Experiment 40. Reaction between iron and sulphur. Have 

ready a Bunsen burner, two test tubes, test tube holder, 
spatula, ring stand with clamp, and magnifying glass; also 
some fine iron filings, iron wire, and flowers of sulphur. 

Take about 5 c.c. of very fine iron filings, and an equal 
volume of flowers of sulphur. Mix them well, either by 
means of a spatula on a piece of paper or by pouring them 
several times from one dry test tube to another. When the 
mixture is of a uniform color, examine it under the mag- 
nifying glass and note the separate particles of both fac- 
tors. Put the mixture in a dry test tube and heat till the 
mixture begins to glow; then take the test tube out of the 
flame and watch the glow increase and spread of its own 
accord thruout the whole mass. When this point has been 
reached, let the test tube cool a little; then empty out the 
residue, or, if it sticks to the glass, break the test tube. Ex- 
amine the residue with the magnifying glass and compare 
it with its appearance before heating. Do the iron filings 
seem to remain unchanged, or does it look as if there had 
been some reaction? Get all the properties you can of the 
resulting substance. Under these conditions it may not be 
possible to change all the iron to iron sulphide, but some 
of it does change. 

Do the next part of this experiment in the hood. 

Clamp a test tube about one fourth full of sulphur to 



FIRST YEAR CHEMISTRY 



109 



the ring stand as in Experiment 37 and heat till the vapors 
of sulphur may be lighted at the mouth of the test tube. 
Heat an iron wire red hot and plunge it into the burning 
sulphur vapors. When a globule has formed on the end of 
the wire, break it open and get its properties. 

Take some iron sulphide from the bottle on the shelf and 
get the properties of iron sulphide from that sample. Write 




Fig. 59. Apparatus for passing hydrogen into warm sulphur vapors. 

the equation for the union between iron and sulphur. We 
have found in many of the experiments we have clone this 
year that oxygen unites with many simple substances to 
form oxides. In the last four experiments we found that 
sulphur unites with several simple substances to form sul- 
phides. This introduces us to the name of a second kind 
of chemical change, namely, synthesis. Bi/ synthesis we 
mean the uniting of two or more simple substances to form a 



110 FIRST YEAR CHEMISTRY 

compound substance. We see, then, that oxidation is only a 
limited and special kind of synthesis. 

Experiment 41. Reaction between hydrogen and warm 
sulphur. (Two students should work together on this ex- 
periment.) Have ready the hydrogen generator with cap- 
illary tube, a 6 by 1 test tube with good cork to fit it, two 
Bunsen burners, two ring stands with rings and clamps, 
tripod, iron rod, round file, bat-wing burner, beaker, grad- 
uate, some soft glass tubing, and some rubber connectors; 
also some roll sulphur, mossy zinc, and sulphuric acid. 

The object of this experiment is to pass hydrogen gas in- 
to hot sulphur vapors, and to examine the product formed. 

Set up the apparatus as shown in Fig. 59. Fit the 6 by 1 
test tube A with a good cork, having two holes in it; thru 
one hole pass a short right-angled glass tube B, extending 
two thirds of the way down into the test tube; connect 
this tube with the outlet tube C from the generator. Thru 
the other hole of the cork in the test tube pass a short right- 
angled tube D, flush with the inner end of the cork. To 
the tube D connect the hard glass capillary tube E by means 
of a rubber connector stout enough to hold the capillary 
tube in horizontal position without further support; the 
capillary tube used in the hydrogen experiments will do 
here, if it is clean; if it is not clean, it is worth while to make 
a new one. Clamp the test tube at such a hight that it 
can be heated with a Bunsen burner, and support the gen- 
erator at an appropriate hight on a ring of the ring stand. 

Study of the new gas. — When the apparatus is properly 
set up, put about 10 grams of roll sulphur in the 6 by 1 
test tube and 10 grams of mossy zinc in the flask. Make 
in a beaker a mixture of 50 c.c. of water and 10 c.c. of con- 
centrated sulphuric acid. Add enough of this 1:5 acid so- 
lution to the flask to seal the thistle tube and give a steady 



FIRST YEAR CHEMISTRY 111 

stream of hydrogen. Light the hydrogen at the capillary 
tip by means of the safety tube. When the hydrogen is 
burning, heat the sulphur till the yellow vapors nearly fill 
the test tube but do not let any of them run over into the 
capillary tube. If the gas is still burning at the tip, blow 
out the flame and note the odor of the new gas. Let the 
reaction run for several minutes till you are sure the new 
gas is being formed. Note its state, color, transparency, 
and odor; other properties may be studied to better ad- 
vantage in the next experiment when the gas is made in 
larger quantities. This new gas is called hydrogen sul- 
phide or sulphide of hydrogen ; occasionally it is called sul- 
phuretted hydrogen. Of what two simple substances must 
hydrogen sulphide be composed? Place the Bunsen burner 
flame under the large part of the hard glass tube near where 
it narrows to the capillary part, and heat a few minutes. 
Continue heating both test tube and hard glass tube till a 
deposit appears in the capillary tube, being careful as be- 
fore that no sulphur vapors pass from the test tube over 
into the capillary tube. What is the deposit? From what 
substance was the deposit in the capillary tube derived? 
What, then, must be going off? Smell of the gas issuing 
from the capillary tube to see if it still has the disagreeable 
odor of the hydrogen sulphide, or if this odor has disap- 
peared. Also light the gas issuing from the capillary tube. 
Is your answer to the last question correct? What is the 
effect of heat on hydrogen sulphide? Write the equation 
for uniting hydrogen and sulphur; in this equation do not 
show zinc and sulphuric acid as factors, for that reaction 
took place in the generator; show simply the reaction that 
took place in the test tube. Also write the equation for 
the decomposition of hydrogen sulphide by means of heat. 
What name did you give in Experiment 17 to the change 
taking place when red oxide of mercury was broken down 



112 FIRST YEAR CHEMISTRY 

by means of heat ? What name might you give to the change 
taking place in this experiment when hydrogen sulphide 
is broken down by means of heat? 

Experiment 42. Action of sulphuric acid on iron sulphide. 
Have ready the generator, four dry fruit jars with washers 
and covers, beaker, large funnel, filter paper, ring stand 
with large ring, gauze, Bunsen burner, graduate, and glass stir- 
ring rod; also some sulphide of iron and some sulphuric acid. 

Generation of gas. — Put ahout 25 grams of iron sulphide 
in small pieces in the flask. Insert the stopper with the 
thistle tube and exit tube; to this side tube should be at- 
tached a rubber tube 20 cm. long; blow thru the rubber 
tube before attaching it to the glass tube, in order to get 
out any of the sulphur powder that generally remains on 
the inside of the rubber tubing after the vulcanizing of the 
rubber. Make up in the beaker a mixture of 75 c.c. of wa- 
ter and 15 c.c. of concentrated sulphuric acid. If the iron 
sulphide is not of first quality, it may be wise to heat the 
mixture of acid and water in the beaker over a Bunsen 
burner till it is near the boiling point, then add a little of 
it to the generator, — enough to seal the tube and to give 
a steady stream of gas. If the action does not start im- 
mediately, wait a little while for it, but do not heat the 
flask after the acid has been added. Fill four fruit jars full 
of the gas by displacement. How did you catch oxygen 
by displacement? 

Treatment of the residue in the flask. — When you have 
caught all the hydrogen sulphide that you need, remove 
the stopper from the flask, set the flask away in the hood, 
and let the reaction run to completion. When no more 
gas is evolved, filter the contents of the flask, evaporate 
the clear filtrate to about one half its volume and set it 
away to crystallize. When crystals have formed, deter- 



FIRST YEAR CHEMISTRY 113 

mine their composition by referring to other crystals you 
have made and then write the equation for the action of 
sulphuric acid on sulphide of iron. 

Study of the gas. — Use one jar to get the properties of 
the gas, particularly its state, color, transparency, odor, 
and solubility in water. To test for this last propert} T , 
add about 50 c.c. of water to the second jar, snap on the 
cover, and shake for some time. Open the jar and test the 
resulting liquid by test paper and by odor to see if any of 
the gas has dissolved in it. In the third jar try the inflam- 
mability of hydrogen sulphide by touching a lighted match 
to the mouth of the jar immediately after taking off the 
cover. Note the color of the flame and the odor in the jar 
after burning. What gas is there, according to the odor, 
in the jar after burning? Was any deposit formed on the 
inside of the jar? Rub the finger around the inside to col- 
lect some of this deposit. What is it? What simple sub- 
stances are there in iron sulphide? W T hat simple substances 
are there in hydrogen sulphide? When hydrogen alone 
burns what is the product? When sulphur alone burns 
what is the product? When hydrogen sulphide burns what 
are the two products probably? Which product did you 
find as a deposit on the inside of the jar? Sometimes there 
occurs a small deposit of flowers of sulphur on the inside 
of the jar due to incomplete combustion of the hydrogen 
sulphide. Did you get any? Write the equation for the 
burning of hydrogen sulphide, assuming that the hydrogen 
and the sulphur both oxidize completely. If you are un- 
certain about the results from burning the hydrogen sul- 
phide, set fire to the gas in the fourth jar and watch closely 
for the products. 

Experiment 43. Formation of a metallic sulphide by pre- 
cipitation. Have ready the generator, with a short piece 



114 FIRST YEAR CHEMISTRY 

of soft glass tubing attached by a rubber connector to the 
exit tube of the generator, a test tube, Bunsen burner, 
beaker, and a graduate; also some copper sulphate, iron 
sulphide, and sulphuric acid. 

Dissolve a small crystal of copper sulphate in half a 
test tube of water to form a clear solution; heat, if neces- 
sary to aid in dissolving. Put two or three small pieces of 
iron sulphide in the generator and make up a solution of 
25 c.c. of water and 5 c.c. of concentrated sulphuric acid. 
Generate just a little hydrogen sulphide again, and pass the 
gas into the copper sulphate solution by dipping the glass 
delivery tube of the generator down into the solution in the 
test tube. Note the black precipitate of copper sulphide 
thrown down, and compare it with the copper sulphide 
made before. Get as many of its properties as you can. 
Write the equation for making copper sulphide from copper 
sulphate by means of hydrogen sulphide. There are only 
two products; the second one is sulphuric acid, which, of 
course, remains in the solution. Where did the copper in 
the copper sulphide come from? Where did the sulphur 
in the precipitated copper sulphide come from? 

Experiment 44. Action of sulphuric acid on zinc sulphide. 

Have ready a test tube, beaker, graduate, and Bunsen 
burner; also the zinc sulphide from Experiment 39 and some 
sulphuric acid. 

Put some of the zinc sulphide made in Experiment 39 in 
a clean test tube. Add a little dilute sulphuric acid, con- 
sisting of one part of acid and five parts of water, and heat 
gently. Note the evolution of gas. Note its odor. What 
is the gas? Recall the reaction between iron sulphide and 
sulphuric acid; then write the equation for the reaction be- 
tween zinc sulphide and sulphuric acid, assuming that zinc 
sulphate is the other product formed without actually evapo- 






FIRST YEAR CHEMISTRY 115 

rating the solution to get the crystals. If you are inter- 
ested in trying the action of sulphuric acid on the sulphides 
of copper and of mercury, do so after reading the following 
note. 

Note on the action of sulphuric acid on sulphides. — Metallic sul- 
phides generally dissolve in sulphuric acid in the same way that the 
two sulphides already studied dissolve in sulphuric acid; the reac- 
tion does not give hydrogen sulphide satisfactorily in those cases where 
the mixture must be boiled very much; this is because at the high 
temperature needed for the reaction some sulphuric acid is likely to 
be vaporized and the fumes of this are so choking that they cover up 
the odor of the hydrogen sulphide. 

Experiment 45. Carbon, its properties and its allotropic 
forms. Take pieces of hard coal, soft coal, charcoal, coke, 
gas retort carbon, and lump graphite; also take some small 
quantities of powdered graphite and of lampblack; also 
make a little soot by holding a glass tube in the luminous 
flame of either burner. Examine all these forms of carbon, 
and get as many properties as possible of each. It might 
be well to arrange the properties in the form of a table, for 
ease in comparison and for future reference. Compare the 
above kinds of carbon with each other, in regard to color, 
luster, hardness, and form. 

Note on carbon. — Carbon is a very common substance, about one 
quarter of one per cent of the earth's crust consisting of this substance. 
One half of the vegetable and animal kingdoms is composed of carbon; 
in the mineral kingdom it occurs both free and combined with other 
substances in such minerals as limestone, marble, chalk, and coral. 
It exists in three allotropic forms, (a) diamond, (b) graphite, and (c) 
the amorphous variety; the last comprises most of the common forms 
of carbon. Diamond is the purest form of carbon; sometimes it is 
found native in crystalline (octrahedral) form; generally, however, it 
is found in irregular, rough fragments which are cut and polished 
to the familiar form seen in rings and other ornaments. Graphite 
(sometimes called plumbago or black lead) is a soft, smutty form 



116 FIRST YEAR CHEMISTRY 

of carbon, sometimes crystalline and sometimes amorphous. Charcoal 
is the black residue left when wood is heated without access 
of air in order to drive out the inclosed gases; it often shows the 
grain and formation of the wood. Animal charcoal, or bone black, 
is made by heating bones to a high temperature. Lampblack is the 
fine, black powder obtained when pitch or tar is incompletely burned. 
Soot is unburned carbon from lamps and chimneys. Coal is made up 
of the remains of the luxuriant vegetation of the Carboniferous Age 
buried under heavy layers of rock and sand. If the layer is found 
near the surface, the mass is porous and the gases have not been pressed 
out; it is then called bituminous coal or soft coal. If the layer has 
been buried deeply in the ground, the gases have been pressed out and 
collected in pockets as natural gas, and the layer is hard and compact ; 
it is then called anthracite coal or hard coal. A poor variety of coal 
that has not been in the ground as long as the other two varieties is 
called lignite. Peat is half decomposed vegetable matter collected in 
swampy places. Muck is undecomposed peat. When soft coal is 
heated in air-tight iron retorts the gases are expelled, gray-black, 
porous, shiny coke is left on the bottom of the retort, and hard, gray, 
shiny, fine-grained gas retort carbon is deposited on the inside of the 
top of the retort. When the gas is passed thru water tar is deposited 
and the cleaned gas is stored in gas-holders for illuminating purposes. 
Many of the substances just mentioned are very impure carbon, some 
of them being less than half carbon. The specific gravity of carbon 
varies with the allotropic form; for diamond it is 3.5, for graphite 2.2, 
and for the other forms generally less than 2 It has never been melt- 
ed or vaporized. The uses of carbon are many. Diamonds are used 
for ornaments, for cutting glass, and for drilling rocks. Graphite is 
used in making "lead" pencils, stove polish, lubricants, and crucibles. 
Many of the amorphous forms are used as fuel for heat and power. 
Animal charcoal is used in filters. Lampblack is an ingredient of 
printer's ink and black paints. Gas retort carbon is made into elec- 
tric light pencils and battery plates. 



Experiment 46. Heating carbon in contact with oxygen. 

First prepare a couple of jars of dry oxygen gas from potas- 
sium chlorate and black oxide of manganese, catching this 
gas by displacement, then bave ready the two jars of oxy- 
gen just made, deflagrating spoon, Bunsen burner, pneu- 



FIRST YEAR CHEMISTRY 117 

matic trough, a rubber tube about 30 cm. long, two small 
beakers, gauze, asbestos sheet, tripod, and some litmus 
paper; also some charcoal; also get a bag of carbon di- 
oxide and a bottle of carbonic acid from the instructor. 

Preparation of carbon oxide. — Make sure the deflagrating 
spoon is clean and dry; then put a little heap of powdered 
charcoal on the spoon, heat it over the Bunsen flame till 
the charcoal begins to glow, and plunge it into a jar of oxy- 
gen, sealing the jar quickly. Compare the burning of car- 
.bon in oxygen with its burning in air. In which did it 
burn more vigorously? Were any dense clouds of fumes 
found in the jar? Was any large amount of powder notice- 
able in the jar, — in addition, of course, to what powdered 
charcoal remained unburned on the spoon? What, then, 
is the state and color of carbon oxide? Loosen the cover of 
the jar, noting whether or not any vacuum has formed. 
Plunge a burning splinter of wood into the jar and note 
what happens to it. Does it burn about the same as it 
does in air, or does it burn more brightly, or does it go out 
promptly? A burning match will not answer in this test, 
because matches are often tipped with sulphur, and it is 
burning wood and not sulphur that we wish to plunge into 
the new gas. Remove the spoon carefully from the jar, 
being careful not to let any of the gas escape. Add about 
50 c.c. of distilled water to the jar, snap on the cover, and 
shake well. Open the jar and test again with a burning 
splinter. Did all the gas dissolve? Test the liquid in the 
jar to see if it is acid. The resulting substance is a very 
weak acid; therefore, use care in testing with litmus paper 
that no other acids are around; also see that the litmus 
paper is distinctly blue, and that it has not been turned 
pink from contact with acid fumes in the air of the labora- 
tory. What was formed when water was added to the first 
oxide of sulphur? What was formed when water was added 



118 FIRST YEAR CHEMISTRY 

to the second oxide of sulphur? What name, then, would 
you give to the compound formed when water was added to 
carbon oxide? Write the equation for the burning of car- 
bon; also write the equation for the formation of carbonic 
acid. 

Study of carbon oxide and of carbonic acid. — In order to 
study some of the other properties of carbon oxide and of 
carbonic acid, get a rubber bag full of carbon oxide, and a 
bottle of carbonic acid from the instructor; always see that 
the carbonic acid bottle is tightly closed when not in use, 
for the acid loses its strength easily, and you need some of 
this acid in one of the following experiments. Invert a 
fruit jar full of water in the pneumatic trough. Attach a 
rubber tube to the nozzle of the rubber bag. Plunge the 
tube from the rubber bag into the trough and let enough 
of the gas escape from the bag to fill the jar. In this sam- 
ple of carbon oxide verify its state, color, odor, and effect 
on a burning splinter; to a jar of it add about 50 c.c. of 
water, shake, and try the action of the resulting liquid on 
test paper. Also test the properties of carbonic acid, us- 
ing some from the bottle. Are the component parts of 
carbonic acid tightly united or not? To answer this ques- 
tion, put about 50 c.c. of carbonic acid in a small beaker 
and an equal amount of distilled water in another small 
beaker; set the two beakers side by side on the wire gauze 
and asbestos sheet, and heat with the Bunsen burner so 
that each beaker gets the same amount of heat. Note in 
which beaker bubbles arise first. Continue heating till the 
distilled water boils and note all differences in the behavior 
of the two liquids, particularly the size of the bubbles and 
the rapidity with which they come. What were the bub- 
bles that arose immediately on the application of heat to 
the beaker containing carbonic acid? What, then, must 
have been left behind in the beaker? Test the remaining 



FIRST YEAR CHEMISTRY 119 

liquid to see if your inference is correct. Write the equa- 
tion for the decomposition of carbonic acid by means of 
heat, showing the two component parts of the acid in full. 

Experiment 47. Action of magnesium on carbon oxide. 

Have ready a dry fruit jar, deflagrating spoon, Bunsen 
burner, and a rubber tube at least 8 or 10 cm. long; also 
a piece of magnesium ribbon about 60 cm. long and a bag 
of carbon oxide. 

Attach the rubber tube to the nozzle of the bag, and plunge 
it well down into the fruit jar. Open the stop-cock of the 
bag just enough to t let a slow stream of carbon oxide enter 
the jar. Hold the cover of the jar lightly over the mouth 
of the jar as the gas is passing in, and from time to time 
bring a burning splinter to the opening at the mouth of the 
jar. When the splinter is extinguished promptly, the jar 
may be considered full of carbon oxide. Withdraw the 
tube from the jar and slip the cover into place. Crumple 
the magnesium ribbon up into a ball, place it on the defla- 
grating spoon, and heat it till the magnesium begins to 
burn, tipping the spoon if necessary to bring the ribbon 
into contact with the flame. When the magnesium is burn- 
ing brightly, put the spoon down into the jar and snap 
on the cover. Note what happens. 

What are the white fumes that fill the jar and settle 
on the inside of the glass? Where did the magnesium 
get the oxygen to form this powder? When opening the 
jar note if any vacuum was formed. What caused this 
vacuum? Break open the ash left on the spoon and recog- 
nize it as free carbon. Where did this carbon come from? 
Write the equation for the action of magnesium on carbon 
oxide. What kind of a chemical change is this? What 
substance was reduced, and to what was it reduced? What 
was the reducing agent? Would you say that the reduc- 



120 



FIRST YEAR CHEMISTRY 



tion was complete? What other substance have you re- 
duced by means of magnesium? 

Experiment 48. Action of hot zinc on carbon oxide. (Two 
students should work together on this experiment.) Have 
ready a piece of hard glass combustion tube 40 cm. long 
and about 1 cm. inside diameter, two rubber bags, each 
supplied with a brass stop-cock, two rubber stoppers to fit 
the combustion tube, two ring stands and two clamps, four 
Bunsen burners, two boxes or other supports for the rubber 
bags, an empty catch bottle, graduate, large beaker, test 




Fig. 60. Apparatus for passing carbon oxide over hot zinc. 

tube, evaporating dish, tripod, gauze, pneumatic trough, a 
piece of rubber tube about 30 cm. long, and some short 
pieces of glass tubing; also some zinc dust and a rubber 
bag of carbon oxide. 

Set up the apparatus as shown in Fig. 60. Connect the 
rubber bags to the combustion tube by means of rubber 
stoppers and short pieces of glass tubing. Clamp the glass 
tube at such a hight that it can be heated by the Bunsen 
burners. Put about 10 grams of zinc dust in the evapo- 



FIRST YEAR CHEMISTRY 121 

rating dish, set it on the gauze over a very low Bunsen flame, 
and heat till the moisture that always gathers in zinc dust 
is all driven out ; be careful not to run the heat high enough 
to oxidize any of the zinc, and use the test tube as a stirring 
rod. breaking apart the lumps of zinc dust with the bot- 
tom of the test tube. Pick out from the zinc dust any 
lumps of zinc that the test tube can not break apart and 
throw them away. When the zinc dust is thoroly dry, 
empty it into the combustion tube and let it spread out 
over the whole length of the tube. Get a bag of oxide 
of carbon from the instructor, making sure that there are 
not over 1000 c.c. of gas in the bag; attach it, with the 
stop-cock closed, to one end of the combustion tube. Let 
the other rubber bag be attached empty to the other end 
of the tube. 

Heat the zinc till it gets quite hot or shows signs of oxi- 
dizing, then open the stop-cock and pass all the oxide of 
carbon to the other bag. Pass the gas slowly back and forth 
from one bag to the other eighteen or twenty times. Be 
sure that you squeeze all the gas out of the bag each time. 
Xote the formation of zinc oxide, — yellow when hot, white 
when cold. Whence came the oxygen to form the zinc 
oxide? Do you think all the oxygen was taken away from 
the carbon oxide? If it were, what would be left? Does 
it look as if any black powdered carbon is left? Has the 
volume of the gas changed to any considerable extent dur- 
ing the experiment? Does the gas that is passed thru 
the combustion tube at the end of the experiment look 
any different from the oxide of carbon that was passed thru 
at the beginning of the experiment? 

Study of the new gas. — Transfer all the resulting gas to 
one bag. Close the stop-cock, remove the bag, attach a 
rubber tube to it, and catch some of the gas in a catch 
bottle over the pneumatic trough. Save the rest of the gas 



122 FIRST YEAR CHEMISTRY 

for the next part of this experiment. Remove the bottle 
from the trough, and touch a flame to the mouth of the bot- 
tle, using a dark background, since the new gas burns with 
only a pale flame. Note the color of the flame. If the 
new gas does not burn, it shows that the carbon oxide had 
not been passed over the zinc long enough. Fill the gradu- 
ate over the pneumatic trough with the new gas, remove 
it from the trough, but keep the hand over the mouth of 
the graduate while you fill a large beaker with water and 
get ready a burning splinter. Then remove the hand, 
touch the flame to the mouth of the graduate, and immedi- 
ately fill the graduate with water from the beaker, pouring it 
in quickly, and pouring it right thru the flame of the burn- 
ing gas. This drives the gas out and enables it to burn 
with a larger flame. What two other gases have you stud- 
ied that burn with a similar flame? Can this new gas be 
either one of those two? Get as many properties of the 
new gas as you can, comparing it especially with the gas 
from which it was made. What property most easily dis- 
tinguishes the two gases from each other? When you re- 
duced oxide of hydrogen by means of magnesium in Experi- 
ment 25, was all the oxygen taken from the water? What 
kind of a change was that? Do you think all the oxygen 
is taken away from the non-combustible oxide of carbon 
in this experiment? Would you still call it reduction? 

Composition of the new gas. — If the non-combustible ox- 
ide of carbon contains only one portion of oxygen, then, 
naturally, when this gas is passed over the zinc and the 
oxygen is removed, carbon would be left. Such, how- 
ever, was not the case. A colorless, combustible gas was 
left. Suppose for a moment that non-combustible oxide 
of carbon contains two portions of oxygen. In that case, 
it might be possible for the zinc to remove one portion of 
oxygen and leave another oxide of carbon less rich in oxy- 



FIRST YEAR CHEMISTRY 123 

gen. Such is the case here. We are justified, therefore, 
in writing the equation for this reduction as follows: — 

f • ) \ carbon | zinc . carbon ) 

< zinc >■ 4- ^ oxygen v = < r + 1 r 

* '( oxy|en ) < ox yg en ) I oxygen J 

(zinc dust) (carbon oxide) (zinc oxide) (carbon oxide) 

To distinguish between the two oxides of carbon on the 
basis of their composition, the combustible oxide of carbon 
is called carbon monoxide and the non-combustible oxide 
of carbon is called carbon dioxide. Write the equation for 
burning carbon monoxide in air, assuming that it is changed 
entirely into carbon dioxide. 

Note on the oxides of carbon. — Carbon monoxide has no common 
name. Carbon dioxide is often found in mines and is then called 
choke damp. In common parlance it is called "carbonic acid" or 
carbonic acid gas. It is not, however, an acid. It is often called 
carbonic anhydride, to show its relation to carbonic acid. There is 
no direct use for carbon monoxide; it is, however, an occasional in- 
gredient of illuminating gas, a regular ingredient of water gas, and 
figures in the reduction of iron ores. Carbon dioxide, on the other 
hand, is used extensively in preparing "soda water," in fire extinguish- 
ers, and in producing very low temperatures. 

Definition of the term "anhydride". — From the statement in the pre- 
ceding paragraph it may be inferred that: An anhydride is any oxide 
that will form an acid when it is chemically united with water. This 
term is related to anhydrous and means waterless. Any particular 
anhydride is designated by the name of the acid from which it may 
be considered as derived by the withdrawing of the water. For in- 
stance, sulphur dioxide is spoken of as sulphurous anhydride, and sul- 
phur trioxide is often called sulphuric anhydride. 

Note on carbonic acid. — Tn Experiment 46 we found that carbon 
oxide united with water to form carbonic acid; we have just now 
found that the ordinary oxide of carbon is carbon dioxide; hence, we 
smut go back to the record of Experiment 46 and correct the equa- 
tions so as to show two portions of oxygen in the oxide of carbon. 
When studying the acids of sulphur we found that the two component 
parts of sulphurous acid were not very tightly united. In sulphuric 



124 



FIRST YEAR CHEMISTRY 



acid, on the other hand, the component parts were tightly united to- 
gether, and we found that it was not necessary to show these compo- 
nent parts when writing sulphuric acid in an equation; no confusion 
arose between the two acids of sulphur thereby, because sulphurous 
acid is an uncommon acid and does not figure in any of our experi- 
ments. The following formulae represent the two sulphur acids, with their 
composition shown in full, then the shortened form for sulphuric acid, 
and lastly the formula for carbonic acid, showing the component 
parts. 



hydrogen \ 
oxygen / 
/sulphur\ 
Voxygen / 



/hydrogen\ 
\oxygen / 
/sulphur\ 
( oxygen I 
Voxygen / 



T hydrogen 
< sulphur 
( oxygen 



(sulphurous acid) (sulphuric acid) (shortened for- 
mula for sul- 
phuric acid) 



hydrogen\ 
oxygen ) 

carbon 

oxygen 

oxygen 

(carbonic acid) 



A shorter form for carbonic acid to correspond with the shorter form 
for sulphuric acid might be written, but it is not advisable to do so. 
On the other hand, whenever carbonic acid figures in an equation it 
should always show its component parts on account of the looseness 
of the union between them. 

Modification of the work on the oxides of carbon. — In Experiment 
48, if you found any marked decrease or increase in the volume of the 
gas as it was changed from the dioxide to the monoxide of carbon, 
such change was probably due either to leakage in the apparatus, or 
to the fact that the temperature of the gas was raised, under which 
latter conditions the volume always increases. Careful heating of the 
combustion tube and accurate measuring of the gases in these experi- 
ments on the oxides of carbon would reveal the fact that one volume 
of carbon dioxide by reduction yields just one volume of carbon mon- 
oxide, and vice versa, one volume of carbon monoxide by oxidation 
yields just one volume of carbon dioxide. It is not necessary to use 
zinc as the reducing agent in Experiment 48; carbon itself would do as 
well. But in that case two volumes of carbon monoxide would result 
from one volume of carbon dioxide. Write the equation for this change, 
i.e., the reduction of carbon dioxide to carbon monoxide by means 
of carbon, and see if you can account for the doubling in volume. 
Then suppose these two volumes of carbon monoxide to be burned in 
air. How many volumes of carbon dioxide would result? 



FIRST YEAR CHEMISTRY 125 

Experiment 49. Chlorine and its properties. Have ready 
two jars of chlorine gas, litmus paper, and turmeric paper. 

Caution. — In testing for the odor of chlorine, breathe 
only the smallest possible amount, because the gas is irri- 
tating to the mucous membrane. The best way to get 
the odor of an unknown gas is to hold the jar a short dis- 
tance before the face and to waft with the hand a little of 
the gas or vapor towards the face. If you happen to in- 
hale too much chlorine, long deep inspirations of fresh, 
cold air is the best antidote, for this washes the chlorine 
out from the air passages. 

In one of the jars get the chief properties of chlorine, 
particularly its state, color, odor, inflammability, solubility 
in water, and action on moist litmus paper and on moist 
turmeric paper. The solubility of chlorine in water may be 
tested as in the case of other gases, — by adding a little 
water to the jar, closing the jar, shaking and testing some 
of the resulting liquid. 

Note on chlorine. — Chlorine itself does not occur free in nature, but 
many of its compounds, called chlorides, do exist in nature in larg e 
quantities. The gas itself must, therefore, be made artificially. Three 
methods ought to be noticed. 

First method. Black oxide of manganese and hydrochloric acid, 
when mixed and heated, give chlorine gas. This method is used ex- 
tensively in the laboratory, but the gas is likely to be impure from 
the fumes of the hydrochloric acid used. 

Second method. Bleaching powder (a complicated compound of 
lime and chlorine) and hydrochloric acid, when mixed in the cold, give 
pure chlorine gas readily. This method is used both in the labora- 
tory and in technical work. 

Third method. If an electric current be passed thru a solution of 
table salt in water, large quantities of chlorine are evolved. This is 
the newest of the three methods and is used largely in technical work. 

Chlorine is twice and a half as heavy as air; this is indicated as 
follows: Specific gravity of chlorine = 2.5 (air = 1). One volume of 
water will dissolve two volumes of chlorine at ordinary temperature. 



126 FIRST YEAR CHEMISTRY 

A solution of chlorine in water is often used instead of the gas itself; 
this solution is called chlorine water. The principal uses of chlorine 
are in the preparation of compounds of chlorine and in bleaching. 

Experiment 50. Reaction between hydrogen and chlorine. 

Have ready the other jar of chlorine gas ordered from the 
storeroom; also a generator with thistle tube and with a 
delivery tube, to which is attached a hard glass capillary 
tip by means of a rubber tube about 30 cm. long, a test tube, 
beaker, and graduate; also some zinc, sulphuric acid, and 
some test papers. 

Generate hydrogen by means of zinc and a 1 : 5 sulphuric 
acid solution. When the hydrogen is coming in a steady 
stream, light it with the safety tube and slowly pass the 
hydrogen flame down into the chlorine. Note the change 
in the color of the flame as chlorine, instead of oxygen, 
joins the hydrogen. Plunge the flame way down to the 
bottom of the jar to make sure that all the chlorine is used 
up, holding the cover lightly in place to prevent escape of 
the gas. Keep the hydrogen flame down in the jar till 
either the flame goes out or it turns back to its original 
color. Note the properties of the product of the union, 
particularly the state, color, odor, action on moist litmus 
paper, and solubility in water. The chemical name of this 
compound is hydrogen chloride, or chloride of hydrogen. 
Compare it carefully with the chlorine that the hydrogen 
chloride was made from, because the odor and action on 
litmus paper are both distinctly different in the two cases, 
tho at first there may seem to be a close resemblance. When 
you have finished this experiment you ought to be able 
to tell by the odor alone whether an unknown gas which 
you may be studying is chlorine or hydrogen chloride. 
Write the equation for the union between hydrogen and 
chlorine, comparing the equation with that for hydrogen 



FIRST YEAR CHEMISTRY 



127 



and sulphur. If you call the uniting of hydrogen and oxy- 
gen "oxidation," what name might you apply to the unit- 
ing of hydrogen with chlorine? 

Note on hydrogen chloride. — The hydrochloric acid that is sold in 
trade is simply a solution of gaseous hydrogen chloride in water. There 
is no union between the gas and the water. The water simply dis- 
solves the gas: it will dissolve four hundred volumes of the gas at 
ordinary temperature. This water solution is more convenient to use 
than the gaseous hydrogen chloride. Whenever it is used, it should 
be expressed in the equation as containing simply hydrogen and chlo- 
rine. A crude variety of hydrochloric acid is sold in trade under the 
name of muriatic acid ; its yellow color is due to impurities. Pure 
hydrochloric acid should be colorless. 



Experiment 51. The action of hydrochloric acid on zinc. 

Have ready an ordinary, large sized test tube with a good 
cork to fit it, the hard glass capillary tip 
used in the preceding experiment, a me- 
dium prescription bottle with a cork to fit 
it, graduate, round file, evaporating dish, 
Bunsen burner, tripod, gauze, and a small 
test tube; also some mossy zinc and some 
hydrochloric acid. 

First make a test tube generator of the 
large test tube by boring a hole thru the 
cork and fitting thereto the glass capillary 
tube as shown in Fig. 61. Next prepare 
a solution of hydrochloric acid for use in 
this and the immediately following ex- 
periments as follows: Put 200 c.c. of dis- 
tilled water in the medium prescription 
bottle; add 40 c.c. of concentrated hydro- 
chloric acid from the bottle on the shelf; 
stopper, and shake to mix well; this makes 
a solution of a strength of 1:5, which is a convenient strength 




Fig. 61. Test tube 
generator. 



128 



FIRST YEAR CHEMISTRY 



for ordinary use. Label the bottle and keep for future 
experiments that part of the solution which is not used here. 
Fill the test tube about two thirds full of the dilute hy- 
drochloric acid just made. Drop in a few pieces of mossy 
zinc, insert the stopper, and light the hydrogen with the 
safety tube. When you are satisfied that it is hydrogen 
that is evolved, remove the stopper, transfer the contents 
of the test tube to a beaker, add some more zinc, and let 
the reaction run, until no more acid is left, heating a little, 
if necessary. How can you tell when this point has been 
reached? Decant, or filter if necessary, and evaporate the 
clear liquid just to dryness in an evaporating dish. Exam- 
ine the residue, which is called zinc chloride, and get its prop- 
erties. Compare what you made with some zinc chloride 
from the bottle on the shelf. Write the equation for the 
action of hydrochloric acid on zinc; omit, of course, in the 
equation the water in which the hydrogen chloride is. dis- 
solved. 



Experiment 52. The action of hydrochloric acid on iron. 

Have ready the test tube generator and the hydrochloric 
acid solution from the preceding experiment, Bunsen burner, 
tripod, gauze, evaporating dish, test tube, beaker, large 
funnel, filter paper, and ring stand with large ring; also 
some iron filings. 

Fill the test tube about one third full of dilute hydro- 
chloric acid of a strength of 1:5. Drop in a few iron filings, 
insert the stopper, and test with the safety tube for hydrogen. 
When you are satisfied that hydrogen is evolved, pour the 
contents of the test tube into a beaker, add a small excess 
of iron filings, and heat the mixture until little or no hy- 
drogen is evolved. Filter, and evaporate the filtrate just 
to dryness in an evaporating dish. Examine the product, 
which is called iron chloride, and get its properties. Com- 



FIRST YEAR CHEMISTRY 129 

pare what you made with some iron chloride from the bottle 
on the shelf. Write the equation for the action of hydro- 
chloric acid on iron. 

Experiment 53. The action of hydrochloric acid on mag- 
nesium. Have ready the test tube generator and hydro- 
chloric acid solution from the preceding experiment, Bun- 
sen burner, tripod, gauze, evaporating dish, and test tube; 
also a piece of magnesium ribbon about 30 cm. long. 

Fill the test tube about one half full with dilute hydro- 
chloric acid. Drop in a piece of magnesium ribbon about 
30 cm. long. Insert the stopper, test for hydrogen with the 
safeti/ tube; then empty the contents into an evaporating dish 
and evaporate to dryness, being very careful not to heat 
more than just to dryness, for high heat decomposes the 
product. Examine the product, which is called magnesium 
chloride, and get its properties. Compare what you made 
with some magnesium chloride from the bottle on the shelf. 
Write the equation for the action of hydrochloric acid on 
magnesium. 

Experiment 54. Sodium and its properties. Get a test 
tube of sodium from the instructor; this element changes 
rapidly when exposed to air, but for this experiment it is 
sufficient protection for the sodium to have it in a small 
stoppered test tube, if it is not kept for any great length 
of time before being used. 

Examine the sodium for its properties, noting particu- 
larly the color, luster, hardness, brittleness, malleability, 
melting point, and inflammability. In testing for the true 
color, be sure to examine a freshly cut surface. Note also 
the permanency of the luster. Omit the test for the solubil- 
ity of sodium in water; sodium reacts with water vigor- 
ously, sometimes with explosive violence, and this reaction 



130 FIRST YEAR CHEMISTRY 

will be studied more fully in one of the following experi- 
ments; therefore, use care here to keep water away from 
the sodium. In testing for the melting point of sodium 
it is sufficient sometimes to knead a little piece of sodium 
between the fingers as you would a piece of putty; if this 
does not reveal anything definite about the melting point, 
then defer the test for this property till you try the effect 
of heat on sodium in the next experiment. 

Note on sodium. — The element, sodium, on account of its activity, 
is not found free in nature. It is made by heating soda and coal to 
a high temperature; the sodium distills off as a vapor and is condensed 
under kerosene or some other oil that does not contain oxygen. An- 
other method recently introduced is to pass an electric current thru 
fused sodium hydroxide, a substance that we shall make in Experi- 
ment 56. The specific gravity of sodium is 0.97; it melts at 95°C, 
and boils at 740°C. Sodium is used in the preparation and purifica- 
tion of alcohol, ether, and similar substances; some is also used in 
making certain compounds of sodium. 

Experiment 55. Heating sodium in contact with air. 

Have ready the cover of the porcelain crucible, pipestem 
triangle, tripod, Bunsen burner, and glass stirring rod; also 
some sodium. 

When you noticed in the preceding experiment that the 
luster of sodium on a freshly cut surface disappeared rapidly 
when sodium was exposed to air, what was happening 
chemically to the sodium? What is the name of the com- 
pound formed in this coating, and also in the white coating 
on the original piece of sodium? Write the equation for 
the spontaneous oxidation of sodium when standing in air. 
Take a small piece of sodium about the size of a pea, put 
it in a clean, dry, and inverted cover of the porcelain cruci- 
ble, and heat very gently. Touch the sodium from time 
to time with a glass rod to find out if the sodium has melted. 
If any scum forms on the sodium, skim it off with the rod, 






FIRST YEAR CHEMISTRY 131 

so as to observe the luster of the melted sodium. Note 
also the permanency of the luster. Increase the heat gradu- 
ally till the sodium catches fire. Note the color of the flame, 
also the color of the residue, which is called sodium oxide. 
Save the resulting sodium oxide for the next experiment. 
If you examined closely the coating formed on the sodi- 
um in the test tube and the residue formed by burning sodi- 
um in air and compared them with each other, you proba- 
bly noticed that tile former was pure white, while the 
latter was yellowish. This difference in color is due to the 
fact that there are two oxides of sodium, — one richer in oxy- 
gen than the other. What two other substances have you 
studied that had two oxides each? The white oxide of 
sodium is the less rich in oxygen of the two oxides and is 
called sodium monoxide ; the yellow oxide of sodium has 
two portions of oxygen and is called sometimes sodium di- 
oxide and sometimes sodium peroxide. The residue left 
after the burning of the sodium probably contains both ox- 
ides, but the dioxide generally predominates. Write the 
equation for the burning of sodium, assuming that only the 
dioxide is formed. 

Experiment 56. Action of water on sodium oxide, or the 
preparation of sodium hydroxide. Have ready the wash- 
bottle, small funnel, filter paper, porcelain evaporating 
dish, small beaker, glass stirring rod, crucible tongs, fruit 
jar, Bunsen burner, tripod, and gauze; also the crucible 
cover containing the sodium oxide from the preceding ex- 
periment, litmus paper, turmeric paper, a small lump of 
steel wool, and a lump of "sodium peroxide." 

Add a few drops of water from the wash-bottle directly 
to the oxide of sodium on the crucible cover as saved from 
the preceding experiment. If the powder does not dis- 
solve readily, add a little more water and stir it. Test the 



132 FIRST YEAR CHEMISTRY 

resulting liquid with blue litmus paper and with turmeric 
paper. Is it acid or is it the opposite of an acid? Com- 
pare with the action of water on other oxides that you have 
made. Substances that have the opposite effect from acids 
on test papers are said to be alkaline. An alkaline substance 
is often called simply an alkali. Empty the liquid on the 
crucible cover into a filter paper in the small funnel and 
rinse off the cover with about 5 to 10 c.c. of water from the 
wash-bottle , letting the wash- water mix with the liquid 
in the funnel; catch the filtrate in the evaporating dish 
and evaporate it to dryness. Examine the new substance 
and record its properties. This substance was formed by 
the direct union of water with sodium monoxide. How 
many component parts are there in this new compound? 
What are they? This new compound is called sodium hy- 
droxide or hydroxide of sodium; sometimes it is called 
sodium hydrate for short. What are the three simple sub- 
stances in sodium hydroxide? Write the equation for the 
union between water and sodium monoxide. Always show 
the two component parts of sodium, hydrate when this sub- 
stance figures in this and in future equations. 

Rub a little of the sodium hydroxide you made between 
the fingers and note the greasy feeling. Leave the dish con- 
taining sodium hydroxide exposed to the air for a few hours. 
What happens to the sodium hydroxide? The moisture 
that it has attracted to itself came from the air, and this 
property is called deliquescence. Deliquescence is the ab- 
sorption of moisture from the air by any chemical substance; 
during the deliquescence the solid is generally changed to 
an aqueous solution. The substance itself is said to be 
deliquescent. What is the name given to the opposite of 
deliquescence, i.e., to water leaving a crystalline substance 
when it is exposed to air ? Mention a number of compounds 
that show this opposite property. 



FIRST YEAR CHEMISTRY 133 

SuJium hydroxide is found in trade in the form of a pow- 
der, but more often in the form of round sticks. Take some 
stick sodium hydroxide from the jar on the shelf and com- 
pare its properties with those of the sodium hydrate you 
made. Put a stick of sodium hydrate in the bottom of a 
small beaker and let it lie exposed to the air for several 
hours and note its deliquescence. 

In the preceding experiment it was found that sodium 
dioxide had two portions of oxygen. Let us try the action 
of water on this substance. Put about 100 c.c. of water 
in the fruit jar and light the Bunsen burner. Have the 
lump of "sodium peroxide" ready in one hand and with 
the other grasp the lump of steel wool lightly with the cruci- 
ble tongs. Drop the lump of sodium peroxide into the wa- 
ter and when the gas begins to be evolved hold the steel 
wool in the Bunsen flame till it glows brightly and then 
plunge the wool down into the jar. What happens to the 
steel wool? What gas is thus shown present? Examine 
the steel wool and note the globules of iron oxide formed, 
comparing them mentally with the iron oxide made some 
time ago. Test the liquid in the jar with litmus paper 
and with turmeric paper. Is it acid or alkaline? What 
compound of sodium is probably dissolved in the water? 
Write the equation for the generation of oxygen from water 
by means of sodium dioxide; also write the equation* for the 
burning of iron in oxygen. 

Note on sodium hydroxide. — Sodium hydroxide is the cheapest of 
several strong alkaline substances on the market; it is sold under the 
name of caustic soda. A solution of sodium hydroxide in water is 
sometimes called soda lye, or sometimes simply lye. Sodium hydrox- 
ide is made on a large scale by boiling soda with lime and evaporat- 
ing the dear liquid to dryness. It is also made by electrolyzing a 
solution of table salt. It i> used extensively in making soap. 

Note on turmeric paper. — In the note on litmus paper at the cud 
of Experiment 29 it was stated that blue litmus is turned red by acids 



134 FIRST YEAR CHEMISTRY 

and back to blue by alkalis. It is customary in some laboratories 
to have red litmus paper in stock for testing for alkalis. A more sat- 
isfactory test paper for alkalis is turmeric paper. Turmeric is also a 
vegetable compound; it is yellow in color, and filter paper, when soaked 
with a solution of turmeric and then dried, is yellow. Alkalis turn 
yellow turmeric paper a deep red brown. Acids of ordinary strength 
have no effect on turmeric. Turmeric paper is, therefore, limited to 
testing for alkalis, but in many laboratories it is preferred to red lit- 
mus paper. 



Experiment 57. The action of sodium on water. Have 

ready a large beaker, large test tube, glass stirring rod, 
evaporating dish, fruit jar, tripod, gauze, Bunsen burner, 
round file, and a piece of coarse copper wire; also some 
sodium, a piece of tin foil, litmus paper, and turmeric paper. 
Fill a large beaker three quarters full of cold water, and 
put in it an inverted test tube, also filled with water. Cut 
off a piece of sodium about the size of two small peas and 
wrap it up in a piece of tin foil. Wrap the foil around the 
sodium in such a way that you have a long, narrow bundle 
that will slip easily into the test tube. Wrap the sodium 
in well, for if it slips out of the foil when under water there 
is an even chance of having an explosion which may throw 
particles of burning sodium around. Then wind a piece of 
copper wire a couple of times around the package to form 
a handle and bend this wire in such a way that when the 
whole is lowered into the water the package can be pushed 
up into the test tube as shown in Fig. 62. With the sharp 
point of the file poke several holes in the tin foil so that the 
water can come into contact with the sodium. Insert the 
sodium in the mouth of the test tube and catch in the in- 
verted test tube the gas evolved. Test the gas by the 
tests for gases that you have already studied. What is 
the gas? Whence must it have come? In that case, what 
became of the oxygen in the water? With what other metal 



FIRST YEAR CHEMISTRY 



135 




Fig. 62. Apparatus for treating 
sodium with water. 



have you reduced water? What were the two products in 
that case? What are the tw r o products in this experiment? 
What did you find in Experiment 56 was the reaction be- 
tween sodium monoxide 
and water? What com- 
pound of sodium does the 
water here contain at the 
end of this experiment? 
What, then, must be the 
action of this solution on 
test paper? Test it and 
see if your inference is cor- 
rect. Dip your fingers into 
the solution and see if it has 
the same greasy feeling 
that the sodium hydrate 
previously m a d e had . 
Filter the solution if any impurities have gotten into it. 
Evaporate a little of the solution to dryness in an evaporat- 
ing dish, and compare it w r ith the sodium hydrate obtained 
in Experiment 56. Write the equation for the reaction be- 
tween sodium and w r ater, assuming that the products are 
sodium monoxide and hydrogen. Then write the equa- 
tion for changing the sodium oxide into the hydrate. Fi- 
nally combine these tw r o equations into one and write the 
equation for sodium on water, showing the final products. 
Remember that it is always best to show the two compo- 
nent parts in sodium hydrate; also take as many portions 
of water as are necessary to get the sodium hydrate. When 
two equations are used, as above, to represent the change 
that takes place they are called twin equations; the use of 
twin equations sometimes helps in making clearer "the in- 
ternal mechanism of the reaction," as it is sometimes called. 
Many, however, prefer the single equation that shows only 



136 FIRST YEAR CHEMISTRY 

the final products. Use whichever form makes the experi- 
ment clearer to you. 

Drop a very small piece of sodium upon the surface of 
some water in a fruit jar or in a large beaker and note what 
happens. Does anything burn? Explain, if you can, why 
the sodium melts and rolls around on the surface of the 
water. 

Some time ago you obtained hydrogen from water by 
means of another metal, but in that case the water was in 
the form of steam. What metal was used in that case? 
What name was given then to pulling oxygen out of an ox- 
ide by means of another substance? Would you consider 
the action of sodium on water the same kind of a reaction? 

Experiment 58. Action of chlorine on sodium. Have 

ready a jar of freshly made and dry chlorine gas; also a 
piece of sodium. 

The sodium should be in the form of a very thin slice, 
freshly cut and about one half of a square centimeter in 
area. Cut the slice of sodium as thin as you can with a 
knife, and then squeeze it out still thinner with the knife 
blade on a sheet of paper. Avoid, as far as possible, letting 
the sodium oxidize. Drop the sodium into the jar of chlo- 
rine, snap on the cover, and set away to stand for several 
days. Examine it from time to time to see if any change 
has taken place in the sodium. When the metallic sodium 
has entirely disappeared, spread the new substance out to 
the air for a short while to let any free chlorine escape from 
it. Get the properties of the powder. What two simple 
substances must there be in the new compound? Write 
the equation for the change. What other substance did 
you unite directly to chlorine? What name did you give 
to that product? What name would you give here to the 
product formed when sodium and chlorine united? Try, 



FIRST YEAR CHEMISTRY 137 

by taste, to recognize the new substance as common table 
salt; do not taste if any unchanged sodium remains. What 
reaction would take place between sodium and the moisture 
on the tongue? The chemical name of table salt is sodium 
chloride, or chloride of sodium. Compare the sodium chlo- 
ride you made with some sodium chloride from the bottle 
on the shelf. 

Experiment 59. Reaction between sodium hydroxide and 
sulphuric acid. Have ready two large beakers, a medium 
beaker, tripod, gauze, Bunsen burner, horn-pan balance, 
set of smaller weights, forceps, graduate, glass stirring rod, 
and test tube; also some stick sodium hydroxide, sulphuric 
acid, crystallized sodium sulphate, litmus paper, and tur- 
meric paper. 

First, prepare a solution of sodium hydroxide for use in 
this and the immediately following experiments. Weigh 
out 20 grams of stick sodium hydrate; dissolve it in 200 c.c. 
of distilled water in a large beaker, heating the mixture, 
if necessary, to hasten the dissolving of the hydrate. Then 
for use in this experiment prepare in another beaker a di- 
lute solution of sulphuric acid by adding 10 c.c. of concen- 
trated sulphuric acid to 100 c.c. of water and stirring or 
shaking to insure complete mixing. 

Pour about 50 c.c. of the recently prepared sodium hy- 
drate solution into a beaker, and add to it about an equal 
volume of the sulphuric acid solution. Stir the mixture 
well with a glass rod having fire-polished ends. Then, 
holding pieces of turmeric paper and of blue litmus paper 
between the fingers of the left hand as shown in Fig. 63, 
touch a drop of the mixture in the beaker by means of the 
stirring rod to each of the two test papers. Note whether 
the mixture is acid or alkaline. If it is acid, add as much 
of the sodium hydrate solution as you think necessary to 



138 



FIRST YEAR CHEMISTRY 



destroy the acidity of the mixture. If, on the other hand, 
the first test shows the mixture to be alkaline, add as much 
of the sulphuric acid solution as you think necessary to de- 
stroy the alkalinity of the 
mixture. Stir well after 
each addition of liquid to 
the mixture before using 
the test paper. Test again 
with test papers. If the 
mixture turns neither test 
paper, it is then neither 
acid nor alkaline, but 
neutral; the process of 
forming a neutral substance 
from an acid and an 
alkali is called neutraliza- 




Fig. 63. The proper way to use 
litmus paper. 



tion. If the mixture still turns one of the test papers, add 
first one and then the other of the two original solutions 
as necessary in small quantities, till, after stirring, a drop 
taken out with the glass rod turns neither test paper. 

Evaporate all of the neutral liquid to crystallization. 
Examine the crystals and get their properties, particularly 
color, form, luster, solubility in water, efflorescence or deli- 
quescence, and water of crystallization. When you treated 
metals with sulphuric acid what name did you give to the 
crystals obtained in each case? When you treated zinc 
oxide with sulphuric acid what did you call the crystals, 
and what became of the hydrogen in that case? In this 
experiment using sodium hydrate and sulphuric acid you 
finally get crystals. What would you naturally call these 
crystals? Have they water of crystallization? Determine 
this by actual experiment. 

Compare the product that you made with some sodium 
sulphate from the bottle on the shelf. What are the two 



FIRST YEAR CHEMISTRY 139 

component parts of crystallized sodium sulphate? Is the 

water of crystallization as tightly united to the anhydrous 
sodium sulphate here as in the case of other crystallized 
sulphates that you have made? Was any hydrogen evolved 
when you neutralized the two solutions? If not, what 
probably became of the hydrogen? If you have any trouble 
in answering this question, see if you have shown the 
component parts in the sodium hydroxide, because this 
helps greatly in keeping track of the various elements in 
the factors and products. Write the equation for the re- 
action between sodium hydrate and sulphuric acid; also 
write the equation for driving water of crystallization out 
of crystallized sodium sulphate. Would the equation for 
the efflorescing of sodium sulphate crystals be the same as 
the last equation? 

Note on sodium sulphate. — Crystallized sodium sulphate is called 
Glauber's Salt in honor of Glauber, the chemist who first made this 
salt. Sodium sulphate is a rather common sulphate, but it is not 
used as extensively as some of the sulphates already studied. 

Experiment 60. Reaction between sodium hydroxide and 
hydrochloric acid. Have ready a large beaker, graduate, 
glass stirring rod, tripod, gauze, Bunsen burner, and test 
tube; also the sodium hydroxide solution left from the pre- 
ceding experiment, hydrochloric acid, litmus paper, and 
turmeric paper. 

Put about 50 c.c. of your prepared sodium hydrate solu- 
tion in a beaker. Then prepare a dilute solution of hydro- 
chloric acid by adding 10 c.c. of concentrated hydrochloric 
acid to 100 c.c. of water and shaking or stirring to mix 
thoroly. To the 50 c.c. of sodium hydrate solution add 
about an equal volume of the dilute hydrochloric acid solu- 
tion. Test the resulting mixture with test paper as in the 
neutralization of sodium hydrate with sulphuric acid. Add 



140 FIRST YEAR CHEMISTRY 

first one and then the other of the sodium hydrate and of 
the dilute hydrochloric acid solutions till the mixture, after 
sufficient stirring, is neutral, i.e., neutralize sodium hydrate 
with hydrochloric acid. 

Evaporate all of the neutral liquid till crystals begin to 
form. Examine the crystals for their properties and rec- 
ognize them as crystallized table salt. Compare what you 
have just made with some sodium chloride from the bottle 
on the shelf. Dry carefully some of the crystals you made 
and heat them in a dry test tube to see if they have any 
water of crystallization. How many factors were there in 
this change? Which factor had component parts? What 
were they? Of what two simple substances is hydro- 
chloric acid composed? What two simple substances are 
there in sodium chloride? Was there any hydrogen evolved 
during this neutralization? What became of the hydrogen 
in the neutralization in the preceding experiment? What, 
then, becomes of the hydrogen in this experiment? How 
many products were there, then, in this experiment ? Write 
the equation for the neutralization of sodium hydrate and 
hydrochloric acid. What one product was common to both 
cases of neutralization that you have just studied? Watch 
all future cases of neutralization to see if this product is 
common to them also. 

Note on sodium chloride. — As already mentioned, sodium chloride is 
commonly called table salt; this name, of course, is derived from the 
use indicated in the name. It is also called sea salt, on account of 
the extensive occurrence of this compound in the ocean. A coarsely 
crystalline variety is called rock salt. Occasionally sodium chloride 
is called simply salt. When the term "salt" is used in every day life 
it refers naturally to sodium chloride; it is well to remember, how- 
ever, that altho the word "salt" means sodium chloride in the labora- 
tory, too, the term "a salt" denotes any compound that may be made 
from an acid by replacing the hydrogen thereof by a metal. To a 
chemist, therefore, a salt may be not only a chloride, but a sulphide, 






FIRST YEAR CHEMISTRY 141 

a sulphate, or one of a number of similar compounds. For the defini- 
tion of the term "salt" refer to the note at the end of Experiment 34. 
Considerable sea salt is obtained by evaporating sea water; some is 
mined. Enormous quantities of sodium chloride are used in making 
other sodium compounds and in the preparation of hydrochloric acid 
and bleaching powder. 

Experiment 61. Reaction between sodium hydroxide and 
carbonic acid. Have ready a large beaker, glass stirring 
rod, tripod, gauze, Bunsen burner, and. test tube; also the 
bottle of carbonic acid, the sodium hydrate solut'on pre- 
pared in Experiment 59, litmus paper, turmeric paper, some 
crystallized sodium carbonate, and some anhydrous sodium 
carbonate. 

Remember that carbonic acid is a weak acid and that 
sodium hydrate is a strong alkali. For this reason, and for 
another that will appear later in the experiment, it is not 
convenient to neutralize sodium hydrate with carbonic acid. 
Instead, take from 75 to 100 c.c. of carbonic acid in a beaker, 
and add some of your prepared sodium hydrate solution to 
it until the mixture after stirring is no longer acid, but 
turns turmeric paper slightly red. If the carbonic acid has 
already lost its strength, get a fresh bottle. 

Evaporate the solution in the beaker to dryness and get 
the properties of the residue. What name did you give to 
the sodium compound resulting from the reaction of sulphu- 
ric acid with sodium hydrate. What name would you give 
to the white residue left in this experiment? What three 
simple substances are there in this white residue? Where 
did the metal in this salt come from? Was any hydrogen 
evolved in this case? What, then, probably became of the 
hydrogen? Write the equation for the reaction between 
sodium hydrate and carbonic acid. 

To get the properties of crystallized sodium carbonate ex- 
amine some from the bottle on the shelf. Select a good, 



142 FIRST YEAR CHEMISTRY 

clear crystal of sodium carbonate and leave it exposed to 
the air for an hour or more. Is it efflorescent or deliques- 
cent? Has the crystallized compound water of crystalliza- 
tion? See if you can find the water of crystallization by 
heating the crystals in the test tube. Compare the residual 
anhydrous sodium carbonate with some of this salt from 
the bottle on the shelf. Write the equation for driving 
water of crystallization out of crystallized sodium carbonate. 
Dissolve a little of either kind of sodium carbonate in 
water and try the action of this solution on turmeric paper. 
Most salts are neutral. It is natural, however, that sodium 
carbonate should have an alkaline reaction, because sodi- 
um forms such a strongly alkaline hydrate, while carbonic 
acid is one of the weakest acids we have studied. 

Note on sodium carbonate. — Crystallized sodium carbonate is often 
called sal soda, sometimes washing soda, and occasionally simply soda. 
Anhydrous sodium carbonate is generally called dry soda to distin- 
guish it from sal soda. Sodium carbonate is one of the most impor- 
tant compounds of sodium, partly because it is a mild alkali, and partly 
because it is the cheapest soluble carbonate known. Furthermore, 
it is used as a reagent in a very large number of reactions. Prepara- 
tion. — On account of its importance, and the fact that it does not 
occur free in large quantities in nature, the technical preparation of 
this salt ought to be considered. Two processes are used, the Le 
Blanc Process and the Solvay Process. The LeBlanc Process con- 
sists of two steps: First, sodium chloride is changed to sodium sul- 
phate by means of sulphuric acid; then the sodium sulphate is heated 
to a high temperature with coal and marble. The resulting mixture 
is treated with water, which dissolves out the sodium carbonate, and 
from this solution the crystallized salt is obtained by evaporation. 
The Solvay Process consists essentially of passing ammonia gas and 
carbon dioxide into a cold, concentrated solution of table salt. This 
process is sometimes called the ammonia process. Enormous quantities 
of sodium carbonate are used in the manufacture of glass and soap. 

Experiment 62. Reaction between sodium hydroxide and 
carbon dioxide. Have ready a large test tube, a glass tube 






FIRST YEAR CHEMISTRY 143 

about 20 cm. long, rubber connector, tripod, gauze, Bunsen 
burner, and evaporating dish; also the sodium hydrate so- 
lution prepared in Experiment 59, a bag of carbon dioxide, 
litmus paper, and turmeric paper. 

Put about 10 c.c. of sodium hydrate solution in a large 
test tube. Get a bag full of carbon dioxide from the in- 
structor, and connect to it by means of the rubber con- 
nector a piece of glass tubing about 20 cm. long. Open 
the stop-cock on the tube of the bag slightly and allow 
all the gas in the bag to bubble slowly thru the liquid in 
the test tube. Evaporate the resulting solution to dryness. 
Examine the residue and secognize it, if possible, as anhy- 
drous sodium carbonate. Try the action of the residue on 
moist test papers. Write the equation for the change that 
took place in the test tube and explain fully how it is that 
you can get a salt in this case when you do not start with 
an acid but with an acid anhydride. 

Experiment 63. Reaction between sodium carbonate and 
sulphuric acid. Have ready two large beakers, a medium 
beaker, graduate, horn-pan balance, set of smaller weights, 
forceps, wooden toothpicks, glass stirring rod, tripod, gauze, 
and Bunsen burner; also some sodium carbonate and some 
sulphuric acid. 

First, prepare for this experiment, and for the one im- 
mediately following it, a solution of sodium carbonate as 
follows: Dissolve 20 grams of sodium carbonate from the 
bottle on the shelf in 200 c.c. of water; either kind of sodi- 
um carbonate will answer the purpose. Put 50 c.c. of this 
solution in a beaker and add to it a dilute solution of sul- 
phuric acid prepared by adding 10 c.c. of concentrated 
acid to 100 c.c. of water. Continue adding the dilute acid 
as long as there is any effervescence. What is the differ- 
ence between effervescence and efflorescence? What is the 



144 FIRST YEAR CHEMISTRY 

gas evolved? Test it by plunging first a glowing and then 
a burning splinter into the clouds of gas that arise from 
the liquid while there is a vigorous effervescence n the 
beaker. Which factor did this gas come from? Since you 
started with sulphuric acid, what salt of sodium have you 
probably in the solution? To determine this definitely, 
evaporate the solution to crystallization and examine the 
crystals, comparing them with the crystals of other sodium 
compounds already studied. Write the equation for the 
change that took place during the effervescence, being sure 
to show the two component parts of sodium carbonate in 
full; the formula for sulphuric acid, however, may be writ- 
ten as in previous experiments, with all the oxygens com- 
bined in one; the formula of the sulphate must, of course, 
correspond to that of the sulphuric acid and that of the car- 
bonate to that of carbonic acid. How many factors were 
there in this experiment? How many products were there? 
This experiment shows a good way to test for a carbonate, 
because sulphuric acid when put upon any carbonate will 
drive off carbon dioxide in considerable quantities, and this 
gas is easily identified. 

Experiment 64. Reaction between sodium carbonate and 
hydrochloric acid. Have ready two large beakers, a medi- 
um beaker, glass stirring rod, graduate, tripod, gauze, Bun- 
sen burner, and several wooden toothpicks; also the sodium 
carbonate solution prepared in the preced'ng experiment, 
and some hydrochloric acid. 

Put 50 c.c. of the solution of sodium carbonate prepared 
in the preceding experiment in a beaker. Then prepare a 
dilute solution of hydrochloric acid by adding 10 c.c. of 
concentrated hydrochloric acid to 100 c.c. of water. Add 
some of this dilute hydrochloric acid to the sodium car- 
bonate solution as long as there is any effervescence. Test 



FIRST YEAR CHEMISTRY 145 

the gas evolved. What is it? Which factor did this gas 
come from? What salt of sodium is probably left in solu- 
tion? Evaporate the resulting liquid till crystals begin to 
form. Get the properties of the resulting crystals, com- 
paring them with crystals of sodium compounds previously 
made. Write the equation for the reaction between sodi- 
um carbonate and hydrochloric acid, remembering, here 
as well as in the preceding experiment, to show the com- 
ponent parts of sodium carbonate; also remember to omit 
water from the formula for hydrochloric acid. What gas 
may be made on a large scale by means of the reaction 
studied in this experiment and in the experiment just 
preceding it ? How many factors were there in this ex- 
periment? How many products were there? In test- 
ing an unknown substance for a carbonate is it necessary 
to use sulphuric acid? What other acid may be used 
instead ? 

Experiment 65. Reaction between sodium chloride and 
sulphuric acid. Have ready the generator used in prepar- 
ing hydrogen, two dry pint fruit jars, beaker, graduate, 
horn-pan balance, set of smaller weights, forceps, Bunsen 
burner, tripod, gauze, and rubber tube about 30 cm. long; 
also some sodium chloride, sulphuric acid, and blue litmus 
paper. 

Put 25 grams of sodium chloride in the generator flask. 
See that the thistle tube nearly touches the bottom of the 
flask. Thru the thistle tube add a little of a solution pre- 
pared by mixing 50 c.c. of sulphuric acid with 25 c.c. of 
water. Do not add the acid solution all at once, but add 
first enough to seal the thistle tube, and then small quanti- 
ties gradually as the gas is evolved. Do not add all of the 
acid solution; use only as much as is necessary to keep the 
tube sealed. If the gas does not come off readily, heat the 



146 FIRST YEAR CHEMISTRY 

flask a little to start the action. Fill the two jars with the 
gas, catching it by displacement of air. Recognize the gas 
by its odor and its action on moist test paper. What is it? 
How have you made this gas before? Be sure that the jars 
are full of hydrochloric acid gas. When no more hydrochloric 
acid gas is evolved from the mixture, even when gentle 
heat is applied to the flask, empty out the liquid residue 
from the flask into a beaker. If the residue is not liquid, 
add cautiously about 50 c.c. of hot water to dissolve the 
residue, heating a little if necessary. Evaporate the clear 
liquid to crystallization and determine the composition of 
the crystals by a study of their properties and by a 
comparison of them with other crystals already made. Write 
the equation for the reaction between salt and sulphuric 
acid. What does this experiment show regarding the rela- 
tive replacing strength of sulphuric and of hydrochloric acids ? 
Compare the equation for the reaction between salt and 
sulphuric acid with that for zinc oxide and sulphuric acid, 
and with that for iron sulphide and sulphuric acid. In 
these three equations we see that the change is not a direct 
combination between simpler substances; nor is it a reduc- 
tion nor an analysis. Instead, there is an interchange be- 
tween the hydrogen of the acid and the metal in the other 
factor, the products being the sulphate of the metal and a 
compound composed of hydrogen and the simple substance 
that was originally in union with the metal. Such a change 
is called a metathesis ; sometimes it is called a metathetical 
change; it may be considered a double substitution. By 
metathesis we mean a chemical change taking place between 
two compound substances whereby a simple substance from 
one of the compounds exchanges places with a simple sub- 
stance from the other compound, the change yielding two new 
compound substances, both of which are different from the two 
original compounds. 



FIRST YEAR CHEMISTRY 147 

This experiment shows us how to prepare hydrochloric 
acid gas on a large scale, and this method is used both in 
the laboratory and in technical work. 

From one jar get any properties of hydrogen chloride 
that you failed to get when you made this substance in Ex- 
periment 50. In the other jar test the solubility of hydro- 
chloric acid gas by adding 25 or 30 c.c. of water to the jar, 
snapping on the cover, shaking, and testing the resulting 
liquid with test papers. 

Experiment 66. Reaction between sodium and mercury. 
Have ready a cover of a porcelain crucible, ring stand with 
clamp, Bunsen burner, two test tubes, the small generator 
used in Experiment 51, large beaker, bat-wing burner, and 
several short pieces of glass tubing; also some mercury, 
sodium, and sodium amalgam from the bottle on the shelf. 

Clamp a clean and dry test tube to the ring stand at an 
angle of about 45°. Empty a little mercury into the hand, 
enough to make a flat globule as it lies in the palm of the 
hand, not over 1 cm. in diameter. Put the mercury into 
the test tube. Then cut off a piece of fresh sodium about 
the size of a pea and drop this in on top of the mercury. 
Stand at arm's length from the apparatus and pass the 
Bunsen flame two or three times across the bottom of the test 
tube. The reaction usually takes place quickly and vig- 
orously. When it has taken place, heat the mixture just a 
little to make sure that no unused sodium remains. Dis- 
connect the clamp from the stand and, using the clamp as 
a test tube holder, pour the contents of the test tube out 
into the cover of the porcelain crucible. Watch it as it 
cools and hardens. Examine the resulting substance. Does 
it resemble either metal started with? The substance 
formed by the union of mercury and sodium is called sodi- 
um amalgam. Compare the sodium amalgam you made 



148 



FIRST YEAR CHEMISTRY 



with some of the sodium amalgam in the jar on the shelf, 
and note its chief properties. 

Add a few c.c. of water to the sodium amalgam that 
you made and let it stand for several hours, or until a glob- 
ule of metallic mercury appears. The mercury in this case 
seems to dilute the sodium, for water does not react as 
readily upon sodium amalgam as upon metallic sodium. 
What happened when metallic sodium acted on water in 
Experiment 57, and what was the product left in solution 
then? Test the solution left from the sodium amalgam to 
see if you have the same product here. In order to test sat- 
isfactorily the gas evolved, set up an apparatus as shown 

in Fig. 64. Replace the capillary 
tip of the small test tube generator 
by a very short delivery tube, 
bent so that the generator can 
rest against the side of a large 
beaker, and so that an inverted 
test tube filled with water can be 
placed over the end of the delivery 
tube. Fill the generator nearly 
full of water, drop in several 
pieces of sodium amalgam, and 
connect the apparatus. When 
the test tube is full of gas, test 
What is left in the bottom of the 
generator? Write the equation for the reaction between 
sodium amalgam and water. Compare this equation with 
the one for the action of sodium on water. 




Fig. 64. 

Apparatus for catching a 
small quantity of gas. 



the gas. 



What is it? 



Definition of amalgam. — It appeared in the above experiment that 
when sodium and mercury were brought into contact there was formed 
a new substance, whose appearance and properties differed from those 
of either factor. Still, it is not quite right to consider sodium amal- 
gam a true compound, because it is possible, by varying the relative 



FIRST YEAR CHEMISTRY 149 

amounts of the two metals, to form a series of sodium amalgams that 
vary in consistency from the liquid mercury to a hard, gray, lusterless 
mass. The amalgam formed might more correctly be considered an 
intimate mixture of the two metals or possibly a solution of sodium 
in mercury, if one cares to extend the idea of solution to substances 
like metals. In the equation in the above experiment it was allowa- 
ble to write sodium and mercury together in brackets, as if we had a 
compound, because the amalgams are the only substances whose com- 
position is of a variable character. Sodium amalgam is not the only 
amalgam known. Mercury unites in a similar manner with a number 
of metals, for instance, zinc, lead, and gold. Have you met any of 
these amalgams in any of your previous work, either inside or out- 
side of the chemical laboratory? 

Sodium and its compounds. — Considerable time has been devoted to 
the study of the compounds of sodium, because they are all common 
substances, and they are used extensively both in the laboratory and 
in technical work. About 2.5 per cent of the earth's crust is sodium; 
this shows that this substance is widely distributed. Furthermore, 
both sodium itself and its compounds are typical of several other 
metals and their compounds. Two of these, potassium and calcium, 
will be taken up next, but these substances need not be studied so 
much in detail as sodium has been studied, particularly if the similari- 
ty between these metals be constantly borne in mind. 

Experiment 67. Potassium and its properties. Get a test 
tube of potassium from the instructor. On account of the 
rapidity with which this metal oxidizes it cannot be kept 
in contact with air, as sodium was, but it must be kept under 
kerosene or some other liquid that does not contain oxygen. 

To get the properties of potassium, take out the piece 
of potassium from the oil, and soak off as much of the oil 
as possible by means of filter paper. In examining for true 
color and luster it is necessary to look quickly upon the 
freshly cut surface. Omit the test for the solubility of po- 
tassium in water; potassium reacts even more vigorously 
than sodium does, and this reaction will be studied more 
fully in one of the following experiments; therefore, use 
care here to keep water away from potassium. In testing 



150 FIRST YEAR CHEMISTRY 

for the melting point of potassium, if slight kneading be- 
tween the fingers does not reveal anything definite about 
the melting point, then defer the test for this property till you 
try the effect of heat on potassium in the next experiment. 
In all the work with potassium, compare it with sodium 
and see wherein it resembles that metal and wherein it dif- 
fers from it. 

Note on potassium. — Like sodium, the element potassium is not 
found free in nature, but its combined compounds are abundant, 
about 2.5 per cent of the earth's crust being potassium. The prepa- 
ration of potassium resembles closely that of sodium; much is now 
obtained by electrolyzing fused potassium hydroxide. The specific 
gravity of potassium is 0.86; it melts at 62.5°C, and boils at 720°C. 
There is practically no use for metallic potassium. 

Experiment 68. Heating potassium in contact with air. 
Have ready the cover of the porcelain crucible, pipestem 
triangle, tripod, Bunsen burner, and glass stirring rod; also 
some potassium. 

The spontaneous oxidation of potassium in air must have 
been noticed in the preceding experiment. In order to get 
a larger quantity of potassium oxide, put a little piece of 
potassium, from which the oil has been soaked as well as 
possible by means of filter paper, in the inverted cover of 
the porcelain crucible and heat it till the potassium catches 
fire. Note particularly the color of the flame. Get the 
more prominent properties of potassium oxide, comparing 
it with sodium oxide. Save the resulting compound for the 
next experiment. Write the equation for the oxidation of 
potassium. Unlike sodium, potassium forms only one ox- 
ide; this corresponds to sodium monoxide, but it is called 
simply potassium oxide. 

Experiment 69. Action of water on potassium oxide, or 
the preparation of potassium hydroxide. Have ready the 



FIRST YEAR CHEMISTRY 151 

wash-bottle, small funnel, filter paper, porcelain evaporat- 
ing dish, small beaker, glass stirring rod, Bunsen burner, 
tripod, and gauze; also the crucible cover containing the 
potassium oxide from the preceding experiment, litmus 
paper, and turmeric paper. 

To the oxide of potassium from the preceding experiment 
add a few drops of water from the wash-bottle. If the 
powder does not dissolve readily add a little more water 
and stir. Test the resulting liquid with test papers. Is it 
acid or alkaline? How does its alkalinity compare with 
that of sodium hydroxide? If the solution is not clear, 
filter it; in that case empty the liquid on the crucible cover 
into a filter paper in the small funnel, and rinse off the cover 
with about 5 to 10 c.c. of water from the wash-bottle; then 
evaporate the filtrate to dryness. Examine the new sub- 
stance which is called potassium hydroxide or hydroxide of 
potassium; sometimes it is called potassium hydrate for 
short. Record the properties of this new compound and 
compare the sample you made with some of the stick potas- 
sium hydrate from the bottle on the shelf. What is its 
most marked property? Write the equation for the forma- 
tion of potassium hydrate. What three simple substances 
does this compound contain, and what are its component 
parts? Always show the two component parts of potas- 
sium hydrate when this substance figures in this and future 
equations. 

Note on potassium hydroxide. — Next to sodium hydrate, potassium 
hydrate is the cheapest and most common soluble alkaline substance. 
It is made on a large scale by a method similar to that for sodium hy- 
drate, namely, by boiling a solution of potassium carbonate with lime 
and evaporating the clear liquid to dryness. It is also made by elec- 
trolyzing a solution of potassium chloride. The product is sold under 
the name of caustic potash. Sometimes it is called simply potash ; 
this last name, however, does not always indicate potassium hydrox- 
ide, because it is applied indiscriminately to several of the common 



152 FIRST YEAR CHEMISTRY 1 

compounds of potassium. A solution of potassium hydroxide in water 
is sometimes called potash lye, or sometimes simply lye. 

Experiment 70. The action of potassium on water. Have 

ready a fruit jar, porcelain evaporating dish, tripod, gauze, 
Bunsen burner, and glass stirring rod; also some potassium, 
litmus paper, and turmeric paper. 

Pour enough water into a fruit jar to cover the bottom. 
Drop into it a piece of potassium as large as a small pea. 
Note carefully what happens. Note that in this case some- 
thing burns. Note also the color of the flame. In what 
respects is the action of potassium on water different from 
the action of sodium on water? Which has the stronger 
action on water, sodium or potassium? Evaporate some 
of the liquid to dryness and get the more prominent prop- 
erties of the residue. What is it? What is its action on 
moist test paper? Recall the action of sodium on water, 
and the idea of representing this reaction by twin equa- 
tions; then write the equations for the action of potassium 
on water; finally represent the entire change by one equa- 
tion. In Experiment 57 the hydrogen set free was caught 
in a test tube, then ignited, but in this experiment the hy- 
drogen set free by the potassium catches fire immediately. 
Write the equation for the burning of the gas evolved from 
the water by means of the potassium. 

Experiment 71. Reaction between potassium hydroxide 
and sulphuric acid. Have ready two large beakers, a medi- 
um beaker, tripod, gauze, Bunsen burner, horn-pan balance, 
set of smaller weights, forceps, graduate, glass stirring rod, 
and test tube; also some stick potassium hydroxide, sul- 
phuric acid, crystallized potassium sulphate, litmus paper, 
and turmeric paper. 

First, prepare a solution of potassium hydroxide for use 
in this and the two immediately following experiments. 






FIRST YEAR CHEMISTRY 153 

Weigh out 20 grams of stick potassium hydrate; disolve it 
in 200 c.c. of water in a large beaker, heating the mixture, 
if necessary, to hasten the dissolving of the hydrate. Then, 
for use in this experiment, prepare a dilute solution of sul- 
phuric acid by adding 10 c.c. of concentrated sulphuric 
acid to 100 c.c. of water, and stirring or shaking to insure 
complete mixing. 

Pour about 50 c.c. of the recently prepared potassium 
hydrate solution into a beaker, and add to it about an equal 
volume of the sulphuric acid solution. Stir well and test 
the resulting liquid with test papers, using the precautions 
given in Experiment 59 as to the use of test papers. Make 
the solution neutral by adding the necessary acid or alka- 
line solution. 'When the solution is neutral, evaporate to 
crystallization and get the properties of the crystals, par- 
ticularly their color, form, and taste. Examine the sul- 
phate of potassium carefully, as you will be referred to this 
compound in a later experiment. Compare the crystals 
that you made with those from the bottle of potassium sul- 
phate on the shelf. 

Hold a crystal of potassium sulphate in the colorless 
Bunsen flame and note the color imparted by this com- 
pound to the flame, i.e., note the flame coloration of potas- 
sium, as it is called. Then hold a crystal of sodium car- 
bonate in the colorless Bunsen flame and note the flame 
coloration of sodium. The flame coloration of potassium 
is characteristic, i.e., all pure compounds of potassium im- 
part this same peculiar color to the Bunsen flame when 
they are held in it. The same holds true for sodium com- 
pounds. Only eight or nine metals give characteristic 
flame colorations, but in each of these cases it is generally 
true that nearly every compound of the metal in question 
will give the flame coloration characteristic of that metal. 

Heat some of the crystals in a dry test tube and note that 



154 FIRST YEAR CHEMISTRY 

no water of crystallization is present. In what respects 
does potassium sulphate differ from sodium sulphate, and 
in what respects is it similar to it? Do you remember 
any other sulphate that had no water of crystallization? 
Write the equation for the neutralization of potassium hy- 
drate with sulphuric acid, remembering to show in full the 
two component parts of the hydrate. 

Note on potassium sulphate. — The main use of potassium sulphate 
is in the preparation of ordinary alum and of potassium carbonate. 

Experiment 72. Reaction between potassium hydroxide 
and hydrochloric acid. Have ready a large beaker, gradu- 
ate, glass stirring rod, tripod, gauze, Bunsen burner, and 
test tube; also the potassium hydroxide solution left from 
the preceding experiment, hydrochloric acid, litmus paper, 
and turmeric paper. 

Put about 50 c.c. of your prepared potassium hydrate 
solution in a beaker, and then prepare a dilute solution of 
hydrochloric acid by adding 10 c.c. of concentrated hydro- 
chloric acid to 100 c.c. of water. Neutralize the potassium 
hydrate solution with hydrochloric acid. Evaporate all the 
solution to crystallization and get the properties of the prod- 
uct. Compare that which you made with some potassium 
chloride from the bottle on the shelf. Dry carefully some 
of the crystals you made and heat them in a dry test tube 
to see if they have any water of crystallization. Write the 
equation for the neutralization of potassium hydrate with 
hydrochloric acid. What one product did you get in this 
experiment that has been common to all cases of neutrali- 
zation you have studied? 

Experiment 73. Reaction between potassium hydroxide 
and carbonic acid. Have ready a large beaker, glass stir- 
ring rod, tripod, gauze, Bunsen burner, and test tube; also 






FIRST YEAR CHEMISTRY 155 

the bottle of carbonic acid, the potassium hydroxide solu- 
tion prepared in Experiment 71, litmus paper, turmeric paper, 
and some potassium carbonate from the bottle on the shelf. 
Before doing this experiment, re-read the directions for 
making sodium carbonate in Experiment 61, in order to 
refresh your memory on the method of procedure. Then 
proceed as follows: Take about 75 or 100 c.c. of carbonic 
acid in a beaker, and add some of your prepared potassium 
hydrate solution to it, till the mixture after stirring is no 
longer acid but turns turmeric paper slightly red. The re- 
action may then be considered complete. The product 
does not crystallize well. Therefore, evaporate the solu- 
tion to dryness and get the properties of the residue. Is 
it deliquescent or efflorescent? What is the action of its 
water solution on test paper? Of what three simple sub- 
stances is potassium carbonate composed? What are its 
two component parts? Write the equation for the reaction 
between potassium hydroxide and carbonic acid. Compare 
the potassium carbonate you made with that in the bottle 
on the shelf. 

Note on potassium carbonate. — Potassium carbonate is one of the 
most important compounds of potassium. It is often called potash; 
sometimes it is called pearlash. Preparation. — Potassium carbonate 
is made by the methods used for making sodium carbonate. It is 
made from potassium sulphate by the LeBlanc Process, and from the 
chloride by the Solvay Process. Some is made from animal and vege- 
table substances, such as grease, wood, molasses, or cream of tartar; 
when these substances are heated to dryness and ignited, potassium 
carbonate remains in the ashes, from which it may be dissolved out 
by water. Considerable quantities of potassium carbonate were for- 
merly obtained by treating wood ashes with water and evaporating 
the solution to dryness. The main uses of potassium carbonate are 
in the manufacture of glass, soft soap, and other potassium compounds. 

Experiment 74. Action of hot potassium on carbon di- 
oxide. Have icad\' a piece of hard glass tube, 20 cm. long, 



156 



FIRST YEAR CHEMISTRY 



and of at least 7 mm. bore, ring stand with clamp and small 
ring, Bunsen burner, glass stirring rod, small beaker, wash- 
bottle, small funnel, filter paper, test tube, wooden tooth- 
picks, and rubber connector; also a rubber bag containing 
carbon dioxide, some potassium, and hydrochloric acid. 

Set up the apparatus as shown in Fig. 65. Clamp the 
hard glass tube at such a hight that it may be heated by 
the Bunsen burner. Put in it a piece of potassium about 

the size of a pea, 
from which the 
o i 1 has been 
soaked by means 
of filter paper. 
To one end of 
the tube attach 
the bag of carbon 
dioxide; this may 
be obtained from 
the instructor. 
The potassium 
should be at least 
10 cm. from the 
open end of the hard glass tube. Open the stop-cock a 
little and pass the carbon dioxide thru the tube till it puts 
out a match at the open end of the tube. Then heat the 
potassium till it melts, still passing the carbon dioxide over 
it. When the temperature has reached the proper point, 
the reaction between the potassium and the gas takes place 
very quickly. Therefore, watch closely as you heat the 
tube. Note that the potassium burns, and that black parti- 
cles of carbon are left. What is the color of the flame? 
Where did the oxygen for the oxidation of the potassium 
come from? With what two other metals have you re- 
duced carbon dioxide? In which of these two cases was 




Fig. 



65. Apparatus for passing carbon 
dioxide over potassium. 



FIRST YEAR CHEMISTRY 157 

the reduction complete, and in which was it incomplete? 
In this experiment was the reduction complete or incom- 
plete? Why should it be complete in some cases and not 
in all cases? 

Note also that a white powder was formed in the hard 
glass tube. To determine the composition of this powder, 
let the tube cool a little, disconnect it from the rubber bag, 
and then with the wash-bottle wash out the deposit from 
the tube into a small beaker, using as small an amount of 
water as possible. Filter from the carbon, catching the 
filtrate in a test tube. Try the action of the filtrate on test 
papers. Add some hydrochloric acid to the filtrate, and by 
means of a burning toothpick test the gas evolved. Is it 
carbon dioxide? 

In Experiments 63 and 64 you found that when an acid 
was put on sodium carbonate, carbon dioxide was evolved. 
All carbonates when treated with an acid yield carbon di- 
oxide. The presence of carbon dioxide in this experiment, 
therefore, indicates a carbonate, and since potassium was 
the only metal present, the white powder must have been 
potassium carbonate. When writing the equation for the 
change that took place in this experiment, it is best to write 
twin equations, i.e., to write first an equation to show the 
reduction of the carbon dioxide to carbon, assuming that 
the other product is potassium oxide, and then to write an 
equation for changing potassium oxide to potassium carbon- 
ate, remembering that an excess of carbon dioxide was passed 
thru the tube. Write the formula for potassium carbonate 
in the equation so that it shows the two component parts. 
When you are sure that you have written the twin 
equations correctly, combine them both into one equa- 
tion, which shall show the two original factors and the 
two final products of the change that took place in the hard 
glass tube. 



158 FIRST YEAR CHEMISTRY 

Experiment 75. Calcium and its properties. Get a piece 
of calcium from the instructor. This element deteriorates 
too rapidly to allow its being put out on the shelf, but not 
rapidly enough to require its inclosure in a stoppered test 
tube when it is given out to the student. 

Examine this metal for its properties, particularly the 
color, luster, hardness, brittleness, malleability, tenacity, 
and melting point. Omit the test for the solubility of cal- 
cium in water, for this reaction will be studied in detail in 
a future experiment. For the complete list of properties 
refer to Experiment 5 on copper. In all the work with cal- 
cium compare it with sodium and potassium and see where- 
in it resembles these metals and wherein it differs from them. 
Also compare the compounds of calcium that you make 
with the corresponding compounds of sodium and potassium. 

Note on calcium. — The element calcium is not found free in nature, 
but its compounds are very freely distributed. Many minerals con- 
tain calcium, the carbonate and the sulphate of calcium occurring 
most frequently. About 3.5 per cent of the earth's crust is calcium. 
The element may be prepared by electrolyzing fused calcium chloride. 
The specific gravity of calcium is 1.6; it melts at 760°C; its boiling 
point has not been determined. No practical use for calcium has 
yet been discovered, but this may be due to the fact that it is only 
recently that the element has been made in any quantity commercially. 

Experiment 76. Heating calcium in contact with air. 
Have ready the crucible tongs and Bunsen burner; also 
some calcium and calcium oxide. 

Sometimes a white coating collects on calcium when it 
lies exposed to the air for any length of time. Did any 
such coating gather on your sample of calcium? What is 
the chemical name for such a coating? In order to get a 
larger quantity of calcium oxide, take a small piece of cal- 
cium, flatten it out on the anvil until it is about as thick 
as the sheet zinc that you studied earlier in the year. Using 



FIRST YEAR CHEMISTRY 159 

the crucible tongs, hold the piece of calcium in the hottest 
part of the Bunsen flame and watch what happens. If the 
calcium is thin enough, it will catch fire and burn. Note 
the color of the flame and the properties of the resulting 
compound. If the Bunsen flame is not hot enough to set 
fire to the calcium, try the blast lamp flame. Write the 
equation for the burning of calcium. 

It is impossible to put out sodium oxide and potassium 
oxide for the use of the class, because these two substances 
deliquesce so readily. With calcium oxide it is different; 
this substance is fairly stable in air. Examine some of the 
calcium oxide from the jar on the shelf, and note any prop- 
erties you could not get readily from the oxide you made. 
Hold a piece of calcium oxide in the hottest part of the 
flame of the blast lamp and note the dazzling white light, 
which is known as the "lime light." 

Note on calcium oxide. — Calcium oxide usually comes in trade in 
lumps under the name of quicklime, or sometimes simply lime. Oc- 
casionally it is called caustic lime. It is made by heating marble or 
limestone to a high temperature in furnaces called "lime kilns"; the 
reaction that takes place in the lime kiln will be studied in one of the 
following experiments. The residue from the lime kiln is the "lime" 
of trade, and this lime must not be exposed to air for any great length 
of time, because under such circumstances it " loses its strength," i.e., 
it undergoes a change as to its chemical composition; this change will 
be studied in the next experiment. 

Experiment 77. Action of water on calcium oxide, or the 
preparation of calcium hydroxide. Have ready a porcelain 
evaporating dish, wash-bottle, two test tubes, glass stirring 
rod, small unnel, filter paper, ring stand with small ring, 
and Bunsen burner; also some calcium oxide, calcium hy- 
drate, litmus paper, and turmeric paper. 

Put a lump of calcium oxide weighing about 25 grams in 
a porcelain evaporating dish. Be sure to select a good lump 



160 FIRST YEAR CHEMISTRY 

of the oxide, for the powder sometimes found in the jar 
usually does not react readily with water. Using the wash- 
bottle, blow water on the lump as long as the water is ab- 
sorbed, but do not add any more water. The lump of cal- 
cium oxide should simply be soaked with water; it should 
not be allowed to lie in water. If no change takes place im- 
mediately, wait a few minutes; sometimes it does not start 
for five or ten minutes. If, after ten minutes no reaction 
has taken place, throw away the contents of the dish and 
start with a new lump. When the reaction has started, 
watch it, and note all phenomena. Let the reaction run 
to an end. When the lump has entirely fallen to pieces, 
examine the powder, which is called calcium hydroxide or 
hydroxide of calcium; sometimes it is called calcium hy- 
drate for short. Get the chief properties of calcium hy- 
droxide. Test the solubility of calcium hydrate by shaking 
a little of the powder with water vigorously in a test tube, 
filtering into another test tube, and exaporating about 10 
c.c. of the clear filtrate to dryness in the test tube. Try 
the action on test papers of moist calcium hydrate and of 
calcium hydrate solution. Expose a little calcium hydrox- 
ide to air for several hours to see if it is deliquescent or efflo- 
rescent. How does it compare with the hydrates of sodium 
and of potassium? Of what three simple substances is cal- 
cium hydrate composed? How many component parts has 
it? What are they? Write the equation for the union be- 
tween water and calcium oxide. Always show the two 
component parts of calcium hydroxide when this substance 
figures in this and in future equations. 

Note on calcium hydroxide. — Calcium hydrate is called slaked lime 
to distinguish it from quicklime. The process of changing calcium 
oxide to the hydrate, as studied above, is called "slaking"; it may be 
seen on a large scale wherever mortar is made. It follows, then, that 
quicklime might be called "unslaked lime"; and so it is occasionally. 



FIRST YEAR CHEMISTRY 161 

When quicklime is exposed to air for a long time it gradually absorbs 
moisture from the air. and crumbles to a powder which consists largely 
of calcium hydrate; the resulting powder is called "air-slaked lime." 
Calcium hydrate is sometimes called simply "lime," but this is not a 
good term to use. because the name "lime" is applied indiscriminately 
in trade to several of the common compounds of calcium. 

Uses of lime. — The uses of calcium oxide are so interwoven with 
the uses of calcium hydrate that it is difficult to consider the uses of 
these two compounds separately. It is, therefore, allowable to speak 
of the uses of "lime." remembering that trade lime is usually the ox- 
ide of calcium, and that the hydrate is seldom sold in trade except in 
small quantities as a chemically pure compound. The various uses 
of lime fall back, then, upon the oxide, altho this is almost always 
converted into the hydrate during the process, and veiy often it is 
the hydrate that should be considered in the use in question. Lime- 
water is a clear solution of calcium hydrate in water. It is used both 
in the laboratory and in the household. Milk of lime, or white-wash, 
as it is often called, is practically lime-water with a small excess of 
powdered calcium hydrate so that some of it is suspended in the liquid, 
making a thin emulsion. It is used in covering ceilings and the in- 
side of cellar walls; also for covering out-buildings and wherever a 
very cheap paint is wanted. Cream of lime, as the name implies, is 
a thick emulsion of calcium hydrate in water, i.e., it is a thick white- 
wash. Its use is for decorative purposes. Mortar is a pasty mixture 
of calcium hydrate, sand, and water. It is generally made by slaking 
quicklime with water and mixing sand with the resulting wet hydrate. 
This mixture is used extensively by masons for binding together bricks 
and stones. The mixture "sets" or hardens, first, by the drying out 
of most of the water, and then by a gradual change of the calcium 
hydrate into other calcium compounds. Plaster is practically mortar, 
with the addition of hair. It is used in plastering walls and ceilings 
of houses, the purpose of the hair being to hold the thick pasty mor- 
tar together until it has had time to set. Other uses of lime include 
purifying illuminating gas, making sodium hydrate and glass, making 
other calcium compounds, bleaching, disinfecting, and fertilizing. 

Experiment 78. The action of calcium on water. Have 
ready the test tube generator, with capillary tube made in 
Experiment 51, small test tube, Bunsen burner, small fun- 
nel, filter paper, ring stand with small ring, evaporating 
dish, tripocl, and gauze; also a small piece of calcium. 



162 FIRST YEAR CHEMISTRY 

Recall the action of sodium on water in Experiment 57. 
Calcium reduces water much more slowly than does sodi- 
um; for that reason the piece of calcium must be pounded 
out on the anvil to very thin sheets or flakes, so that the 
water can come into contact with a large surface of the metal. 
Fill the test tube about two thirds full of water. Drop into 
it the thin sheets of calcium and quickly insert the stopper, 
which should be fitted with a short piece of glass tubing, 
drawn out to a tip. Note the action going on in the test 
tube. Test the gas evolved with the safety tube. Where did 
the hydrogen come from? What united with the oxygen 
of the water in the place of the hydrogen? Recalling the 
action of water on calcium oxide in Experiment 77, what 
compound of calcium would you say the solution probably 
contains? What, then, must be the action on turmeric 
paper of the liquid in the test tube? Test the liquid to see 
if your inference is correct. If the solution is still clear and 
colorless, and some calcium still remains, let the reaction 
run a little longer, and watch it to see if any white, insolu- 
ble powder is formed. What name would you give to that 
powder? Filter a little of the liquid and evaporate to dry- 
ness. What is the residue from the evaporation? Write 
the equations for the action of calcium on water, using, 
first, the twin equations and then combining them both into 
one. In what two other cases did you use twin equations? 

Experiment 79. Reaction between calcium hydroxide 
and hydrochloric acid. Have ready a 500 c.c. flask, large 
funnel, filter paper, evaporating dish, tripod, gauze, Bunsen 
burner, glass stirring rod, large prescription bottle with 
cork to fit it, and some beakers; also some calcium hydrate, 
hydrochloric acid, litmus paper, and turmeric paper. 

Preparation of lime-water. — First, prepare for use in this 
experiment, and those immediately following it, an aqueous 



FIRST YEAR CHEMISTRY 163 

solution of calcium hydrate. This solution is called lime- 
water and should be made as follows: Take enough finely 
powdered calcium hydrate to fill a small beaker. Trans- 
fer it to the 500 c.c. flask, using the glazed paper method 
described in Experiment 19 and illustrated in Fig. 38; then 
fill the flask from one half to two thirds full of distilled 
water and shake vigorously. This shaking will, of course, 
dissolve a little of the calcium hydrate, and at the same 
time all of those impurities which are soluble. Let the 
mixture settle, then pour off as much of the clear liquid 
as possible, and throw it away. Again, fill the flask about 
two thirds full of distilled water and shake. This time 
nearly pure calcium hydrate dissolves in the water, but the 
solution takes place slowly. Let the unused calcium hy- 
drate settle and then filter the clear liquid directly into the 
large prescription bottle. The filtrate should not be caught 
in a beaker, because when lime-water is exposed to air a 
scum gathers on it, which interferes with some of the later 
experiments where lime-water must be used. The lime- 
water in the prescription bottle will keep for an indefinite 
time, but it must be clear, and it must be tightly stoppered 
when not in use. The powdered calcium hydrate left be- 
hind in the flask may be shaken with fresh portions of dis- 
tilled water as many times as needed to yield the required 
amount of clear lime-water. 

Put about 100 c.c. of lime-water in a beaker and add im- 
mediately small portions of dilute hydrochloric acid of a 
strength of one part of acid to five parts of water, till the 
mixture, after stirring, is slightly acid. Evaporate the so- 
lution to dryness, and get the properties of the residue, 
which is called calcium chloride. Compare your product 
with some of the calcium chloride from the bottle on the 
shelf. Allow some calcium chloride to lie exposed to air 
•veral hours, to find out whether it is deliquescent or 



164 FIRST YEAR CHEMISTRY 

efflorescent. Write the equation for the reaction between 
calcium hydrate and hydrochloric acid, referring to the 
equation for the neutralization of sodium hydrate with 
hydrochloric acid, if you are in doubt. 

Note on calcium chloride. — On account of the ease with which it 
absorbs moisture, calcium chloride is used extensively for drying gases. 

Experiment 80. Preparation and properties of calcium 
sulphate. Have ready several test tubes, Bunsen burner, 
small funnel, filter paper, ring stand with small ring, evapo- 
rating dish, and glass stirring rod; also some calcium chlo- 
ride, sulphuric acid, gypsum, plaster of Paris, powdered 
calcium hydrate, and lime-water. 

Dissolve one small lump of calcium chloride in half a 
test tube of water, heating, if necessary, to aid in the dis- 
solving. To the clear solution of calcium chloride add a 
little dilute sulphuric acid of a strength of 1:5. Recall the 
action of sulphuric acid on sodium chloride. In that case, 
hydrogen chloride was evolved as a gas. In this experi- 
ment, hydrogen chloride, tho formed, is not evolved, be- 
cause there is so much water present, and the water dis- 
solves the hydrogen chloride. What is the other product, 
which falls out as a white precipitate? A precipitate is any 
insoluble substance that separates when two solutions are 
mixed; a precipitate usually falls to the bottom of the test 
tube or other vessel containing the mixed solutions. The 
white precipitate in this case is called calcium sulphate, 
and to indicate the method of preparation it is usually 
called precipitated calcium sulphate. Get the properties of 
calcium sulphate, particularly its form, and solubility in 
water. To test solubility, filter the mixture in the test tube, 
put a little of the precipitate in a clean test tube, add some 
distilled water, and heat. Then filter again, and evaporate 
a little of the clear filtrate to dryness. Was the calcium 



FIRST YEAR CHEMISTRY 165 

sulphate wholly soluble, partially soluble, or insoluble? 
Write the equation for the metathesis between calcium 
chloride and sulphuric acid. 

Note on calcium sulphate. — Calcium sulphate occurs in two forms, 
first, as a white, amorphous powder similar to that made above, and 
second, as a semi-transparent lustrous crystal or crystalline mass. 
The amorphous variety is called plaster of Paris, the crystalline is 
called gypsum. Gypsum occurs in nature as a mineral and contains 
water of crystallization. Plaster of Paris is made by heating gypsum 
to a rather high temperature in order to drive off the water of crystal- 
lization. 

Further study of calcium sulphate. — Put a few small pieces 
of gypsum, i.e., of crystallized calcium sulphate, in a clean 
and dry test tube, and heat till the water of crystallization 
is deposited on the sides of the test tube. Get the proper- 
ties of anhydrous calcium sulphate. Write the equation 
for driving water of crystallization out of crystallized cal- 
cium sulphate. 

Anhydrous calcium sulphate, i.e., plaster of Paris, when 
moistened with water sets up into a hard mass. Take a 
little plaster of Paris from the bottle on the shelf, put it in 
an evaporating dish, and add water until the mass, when 
stirred well, forms a thick cream. Allow it to stand a while, 
and when it begins to harden, press a coin down into the 
stiffening mass. When the plaster of Paris has become 
quite hard, remove the coin and note the impression of it 
in the plaster. The experiment just performed is typical 
of making all kinds of plaster casts. 

The solubility of calcium sulphate is of considerable im- 
portance in elementary chemistry and in analytical work. 
Test the solubility of gypsum in the same way that you 
have tested the solubility of other compounds, i.e., treat a 
few pieces of gypsum with an excess of distilled water in a 
best tube, heating if necessary. Decant a little of the clear 



166 FIRST YEAR CHEMISTRY 

solution into a clean test tube, and evaporate this to dry- 
ness. Is any residue left? What is it? How does the 
solubility of gypsum compare with the solubility of precipi- 
tated calcium sulphate? How does the solubility of cal- 
cium sulphate compare with the solubility of other sulphates 
that you have made? Judging from your work with plas- 
ter of Paris, what would you say about the solubility of 
calcium sulphate? What more definite information does 
the experiment on gypsum just performed give in regard 
to the solubility of calcium sulphate? 

Put some powdered calcium hydrate in an evaporating 
dish, add water until you have a thick paste, and then add 
slowly, with constant stirring, from 5 to 10 c.c. of concen- 
trated sulphuric acid. If nothing happens immediately, 
stir for a couple of minutes and set the resulting mixture 
aside. Recall the neutralization of sodium hydrate with 
sulphuric acid, and then write the equation for the reaction 
between calcium hydrate and sulphuric acid. Note that 
the calcium sulphate sets up pasty or even solid, and can- 
not be distinguished easily from any unused calcium hydrate. 
Would you consider this a good method for making calci- 
um sulphate? Would it improve the method any, if you 
added an excess of sulphuric acid so as to transform all the 
calcium hydrate to the sulphate? 

Finally, put a little clear lime-water into a clean test tube 
and add to it a little dilute sulphuric acid. From the early 
work in this experiment one might naturally expect white, 
amorphous calcium sulphate to be precipitated. Do you 
find it so? How can you account for the fact that no pre- 
cipitate appears? If you add an excess of sulphuric acid, 
the alkaline reaction due to the lime-water disappears. 
Try this and prove that it is so. This indicates, of. course, 
that a reaction has taken place. What light does this ex- 
periment throw upon the solubility of calcium sulphate? 



FIRST YEAR CHEMISTRY 167 

Do you consider the preparation of calcium sulphate from 
lime-water and sulphuric acid a good method for prepar- 
ing calcium sulphate on a large scale? Of all the methods 
we have studied for making calcium sulphate, which one 
do you consider the best one for making calcium sulphate 
on a large scale? 

Experiment 81. Reaction between calcium hydroxide 
and carbonic acid. Have ready a test tube and Bunsen 
burner; also some lime-water and carbonic acid. 

Fill a large test tube about one quarter full of lime-water. 
Add to it a very few drops of carbonic acid. If the carbon- 
ic acid is in good condition, a white precipitate should ap- 
pear. What is this white precipitate? What is the solu- 
bility in water of the precipitated calcium carbonate? Now 
add to the test tube an excess of carbonic acid. What 
happens to the precipitate? What does this show in re- 
gard to the solubility of calcium carbonate in carbonic 
acid? Boil the clear carbonic acid solution in the test tube 
until the free carbonic acid has been destroyed. Note the 
precipitate reforming. Why does the calcium carbonate 
drop out again as a solid when the solution is boiled ? Write 
the equation for the reaction between calcium hydrate and 
carbonic acid, referring to the reaction between sodium 
hydrate and carbonic acid, if in doubt. Also write the 
equation for the decomposition of carbonic acid by means 
of heat. 

Note on calcium carbonate. — Calcium carbonate, like calcium sul- 
phate, occurs in two forms. The first, or amorphous form, of which 
the precipitate made above is an example, is called chalk; sometimes 
it is called precipitated calcium carbonate. The second variety, i.e., 
the crystallized variety, occurs in nature as a number of minerals, 
all of them known under the name of calcite. The more important 
varieties of calcite aro limestone, marble, and Iceland spar ; the first 
two generally show a little crystalline structure; the third variety is 



168 FIRST YEAR CHEMISTRY 

distinctly crystalline. Other varieties of calcium carbonate are called 
stalactites, stalagmites, onyx, and coral. Calcium carbonate is used 
extensively in making glass, in the extraction of iron from its ores, in 
making quicklime, and for building purposes. 

Experiment 82. Reaction between calcium hydroxide 
and carbon dioxide. Have ready two test tubes, Bunsen 
burner, a glass tube about 20 cm. long, and a rubber con- 
nector; also some lime-water and a rubber bag of carbon 
dioxide. 

Fill a clean test tube about two thirds full of lime-water, 
and pass into it some carbon dioxide, using the method ex- 
plained in Experiment 62 for passing carbon dioxide into 
sodium hydroxide solution. Note that when the carbon 
dioxide is first passed in, a white precipitate appears, but 
that this precipitate disappears as more carbon dioxide is 
passed in. What is the white precipitate? Why does it 
first appear, and then disappear? Boil the resulting liquid 
to see if your explanation is correct. Write the equation 
for the reaction between calcium hydrate and carbon di- 
oxide. 

Again, take about 3 or 4 c.c. of lime-water in a test tube 
and by means of a glass tube blow air from the mouth thru 
the liquid. Does a white precipitate appear? What is it? 
What does this show in regard to the presence of carbon di- 
oxide in the breath? Continue blowing thru the liquid to 
see if you can make the precipitated calcium carbonate 
dissolve. 

Experiment 83. Decomposition of calcium carbonate by 
means of heat. (Two students should work together on 
this experiment.) Have ready a piece of hard glass com- 
bustion tube 20 cm. long and about 1 cm. inside diameter, 
two good corks to fit the combustion tube, ring stand with 
clamp, two Bunsen burners, two sets of catch bottles, sue- 



FIRST YEAR CHEMISTRY 



169 



tion pump, wash-bottle, several short pieces of glass tubing, 
and of rubber tubing; also some marble, lime-water, and 
turmeric paper. 

Set up the apparatus as shown in Fig. 66. Clamp th e 
combustion tube at a convenient hight to be heated by th e 




Fig. 66. Apparatus for decomposing calcium carbonate by 
means of heat. 

Bunsen burners; fit each end of it with a good cork, thru 
which passes a short piece of glass tubing. To one of these 
short tubes attach the entrance tube of a single catch bottle; 
connect the exit tube of this catch bottle to the suction 
pump; to the short glass tube at the other end of the com- 
bustion tube attach a set of two catch bottles. 

When the apparatus has all been set up put 5 or 6 small 
lumps of marble in the combustion tube, and fill the three 
catch bottles properly full of clear lime-water. Heat the 
marble with a couple of Bunsen burners for about five 
minutes. Then turn on the water just a little so that air may 
be drawn in thru the double catch bottles, over the heated 
marble and out thru the single catch bottle. This passing 
of air thru the apparatus carries off the gaseous products 
of decomposition of the marble and passes them thru the 



170 FIRST YEAR CHEMISTRY 

lime-water in the single catch bottle. The decomposition 
of the marble does not take place readily, unless the gaseous 
products of decomposition are carried off by means of an 
air current. The oxygen in the air passed over the marble 
does not enter into the reaction. Continue heating the 
marble for 15 minutes, while air is being drawn thru the 
apparatus. 

Ordinary air almost always contains a little carbon di- 
oxide. This impurity of carbon dioxide must, of course, 
be completely washed out of the air before the air is passed 
thru the apparatus, otherwise this trace of carbon dioxide 
might interfere with later results in this experiment. The 
lime-water in the first of the set of two catch bottles will 
probably become milky. That in the second one should 
remain clear; this will indicate that all the carbon dioxide 
in the air is being taken up in the first bottle, and that air 
perfectly free from carbon dioxide is being passed over the 
marble. 

What change occurs in the lime-water in the single catch 
bottle as the air and the gaseous products of the decom- 
position bubble thru it? What gas causes this turbidity 
in lime-water? Whence came the carbon dioxide that re- 
acted with the lime-water in the single catch bottle ? What 
detail of the experiment assures you that your answer to 
this last question is correct? What, then, must be left in 
the combustion tube? Empty out the solid residue from 
the heated marble, and pick out those pieces that show 
greatest evidence of change; put one of the pieces on a 
strip of dry, turmeric paper and add carefully a couple of 
drops of water from the wash-bottle. Is there any change 
in the turmeric paper? What is indicated by this change? 
What does this show regarding the composition of the resi- 
due from the heated marble? Put a lump of the residue 
in the palm of tho hand, and add carefully a couple of drops 



FIRST YEAR CHEMISTRY 171 

of water from the wash-bottle. Is there any evidence that 
the lump is slaking? 

Write the equation for the decompositon of calcium car- 
bonate by means of heat. Is this analysis proximate or 
ultimate? Write the equation for the reaction that takes 
place in the single catch bottle. Also write the equation 
for the reaction that takes place when water is added to the 
residue from the heated marble. 

The main reaction studied in this experiment is the one 
referred to in the "Note on calcium oxide," under Experi- 
ment 76, where the preparation of quicklime from limestone 
is described. 

Experiment 84. Reaction between calcium carbonate 
and hydrochloric acid. Have ready the large generator 
used in preparing hydrogen in Experiment 32, test tube, 
glass tube about 20 cm. long, evaporating dish, tripod, 
gauze, Bunsen burner, graduate, beaker, and rubber con- 
nector; also some marbl^, hydrochloric acid, and lime- 
water. 

Put some small pieces of marble in the generator flask, 
and insert the stopper with the thistle tube and delivery 
tube; then add dilute hydrochloric acid of a strength of 
1:5 thru the thistle tube, and pass the gas evolved thru 
some lime-water in a test tube. What is the gas evolved? 
Whence came it? What is the white precipitate in the 
test tube? What became of the calcium in the marble? 
Evaporate a small quantity of the clear liquid in the flask 
to dryness and recognize the residue as calcium chloride. 
Write the equation for the reaction between calcium car- 
bonate and hydrochloric acid, remembering, as before, to 
show the component parts of the carbonate. How many 
factors are there in this experiment? How many products 
are there? What other carbonate have you treated with 



172 



FIRST YEAR CHEMISTRY 



hydrochloric acid? How many products were there in that 
case? 

This experiment shows a good method for making a cer- 
tain gas on a large scale. What gas? This experiment 
also shows the complete test for a carbonate. This test 
needs to be applied in a number of cases in this year's work, 
so it is well to remember that the best test for a carbonate 
is to treat the supposed carbonate with hydrochloric acid, and 
to pass any gas evolved thru lime-water. The appearance of 
a white precipitate indicates that the substance started with 
was a carbonate. In what respects is this method here de- 
scibed better than the. incomplete test given in Experi- 
ment 63? 



Experiment 85. Reaction between calcium carbonate and 
sulphuric acid. Have ready the large generator with de- 
livery tube, graduate, and test tube; also some marble, sul- 
phuric acid, and lime-water. 

Put a few small pieces of marble in the generator flask, 
insert the stopper with the thistle tube and delivery tube, 
add about 100 c.c. of water, and then from 10 to 20 c.c. 
of concentrated sulphuric acid. As soon as the gas is evolved 
pass it thru some lime-water in a test tube. What hap- 
pens to the lime-water? What does this show in regard 
to the composition of the gas? Whence came the gas? 
What is the white precipitate in the test tube? Note that 
the calcium sulphate which forms one of the other products 
has a tendency to set up hard around the lumps of marble 
and thus prevent continued action. If you should need to 
generate carbon dioxide from marble on a large scale, which 
acid would you consider more convenient to use, hydro- 
chloric or sulphuric? Write the equation for the reaction 
between calcium carbonate and sulphuric acid. How many 
products are there in this case? Also write the equation 



FIRST YEAR CHEMISTRY 173 

for the reaction that takes place in the test tube of lime- 
water. 

Experiment 86. Hard water and soft water. Have ready 
a test tube rack with several test tubes, glass stirring rod, 
and wash-bottle containing distilled water; also some soap 
solution, calcium sulphate solution, lime-water, and car- 
bonic acid. 

Definition of soft and of hard water. — Water that has noth- 
ing in solution is called soft water ; distilled water is a good 
example of soft water; ordinary rain water is usually soft. 
Water that has in solution some compound of calcium is called 
hard water. Hard water is generally divided into two 
kinds, — permanently hard water and temporarily hard 
water. Permanently hard water is water which has dis- 
solved calcium sulphate; its name is derived from the fact 
that it cannot be softened, except by boiling off the water 
as steam and condensing this, leaving the solid calcium 
sulphate behind. Obviously, this is an expensive method. 
Temporarily hard water is water which is charged with car- 
bon dioxide to such an extent that it is able to hold some calci- 
um carbonate in solution by virtue of the solubility of this 
compound in carbonic acid; its name is derived from the 
fact that the water is hard only as long as the carbonic 
acid is present; when temporarily hard water is boiled the 
carbonic acid is destroyed, with consequent precipitation 
of insoluble calcium carbonate; the resulting water is soft. 

Test for hard water. — The difference between hard and 
soft water is shown best by means of soap. Take a little 
soap solution from the bottle on the shelf; this is simply a 
clear solution of pure castile soap in distilled water. Care 
must be taken not to shake the bottle containing the soap 
solution, for a clear soap solution and not soap lather is 
needed for this work. Add two or three drops of the clear 



174 FIRST YEAR CHEMISTRY 

soap solution to half a test tube of distilled water. Mix 
well by means of a glass rod, but do not shake. Is any 
precipitate formed, or does the liquid remain clear? Then 
put the thumb over the mouth of the test tube, shake a 
little, and note the ease with which a lather is produced. 
Next take half a test tube of calcium sulphate solution 
from the bottle on the shelf and add two or three drops 
of soap solution. As before, look for a precipitate before 
shaking and then shake to see if a lather forms as easily as 
in the case of soft water. Finally take about 5 c.c. of lime- 
water in a test tube and add carbonic acid to it till the pre- 
cipitate which is formed at first dissolves. To this tempo- 
rarily hard water add a few drops of soap solution, look for 
a precipitate, shake to get a lather, and compare with soft 
water with respect to ease of getting a lather. The slimy 
precipitate formed when soap solution is added to either 
kind of hard water is called "lime soap." It is formed by 
a metathesis between the soap and the calcium compound 
in the hard water. The sodium from the soap and the cal- 
cium from the hard water change places. Soaps have very 
complicated formulae, so it is not worth while to try to 
write the equation for the metathesis. 

A chemical investigation. — Up to this point in the year's 
work we have spent most of the time making new com- 
pounds and studying their properties. In nearly all cases, 
however, we have made the compounds from simpler sub- 
stances whose composition we knew, or we have made 
them directly from the elementary substances themselves. 
But practical, everyday problems in chemistry are by no 
means as simple as the experiments we have been doing. 
More often a complicated compound, of whose composition 
we are totally ignorant, comes to us, and this compound 
we have to study, either by itself, or in its reaction with 



FIRST YEAR CHEMISTRY 



175 



other substances, the composition of which we know; from 
the data thus obtained, we must be able to tell of what 
the compound is composed. Let us now turn our attention 
to such a problem, and apply what we have already 
learned to the investigation of an unknown substance. 
In warm countries there is often found a white salt-like 
coating on the soil, particularly where animal and vegeta- 
ble matter has been decaying. This deposit is called niter 
or saltpeter, and its composition is different from that of 
compounds so far studied. Since our previous work has 
shown that sulphuric acid reacts with many compounds, 
let us try first the action of sulphuric acid on niter. 



Experiment 87. Reaction between sulphuric acid and 
niter, or the preparation 
of nitric acid. Have 
ready a glass-stoppered 
retort, ring stand with 
clamp, tripod, gauze, 
Bunsen burner, several 
test tubes, test tube 
rack, glass stirring rod, 
horn-pan balance, set of 
smaller weights, gradu- 
ate, large funnel, evapo- 
rating dish, and several 
large beakers; also some 
niter, sulphuric acid, and 
litmus paper. 

Set up the apparatus 
as shown in Fig. 67. 
By means of the clamp 
and ring stand support the glass retort A so that its 
bottom rests lightly on the gauze, and its stem B is in- 




Fig. 67. Apparatus for decomposing 
niter with sulphuric acid. 



176 FIRST YEAR CHEMISTRY 

clined at such an angle that it can extend well down 
into a large test tube immersed in a large beaker of water. 
The test tube should be pushed well down into the water, 
clear to the bottom of the beaker, if possible. See that 
the retort is dry inside and that the glass stopper fits air 
tight. Put 30 grams of niter in the retort, pouring it in 
thru the tubulature, C, using the glazed paper method de- 
scribed in Experiment 19. Be careful not to get any niter 
into the stem of the retort. Then add 10 c.c. of concen- 
trated sulphuric acid, pouring that, too, thru the tubula- 
ture, and using a funnel in order that none of the sulphuric 
acid shall run into the stem of the retort. 

Heat the mixture in the retort for some time. Do not 
hold the hand where the molten mass in the retort can harm 
it, should the retort happen to break. Do not let the tem- 
perature run so high that vapors from the mixture roll out 
thru the stem of the retort. The vapors should condense 
to a liquid in the stem and the liquid should drop slowly 
from the end of the stem into the test tube. Note the 
color of the vapor in the retort at different stages of the ex- 
periment. Note also any changes in the color of the va- 
por. The process of changing a liquid to a vapor and then 
condensing the vapor again is called distillation ; the resulting 
liquid is called the distillate. Continue the heating till you 
have caught 10 or 15 c.c. of the distillate, or till no more 
distillate drops from the end of the stem of the retort. Get 
the properties of this liquid, diluting it with considerable 
water in case the liquid does not react satisfactorily with 
the test papers. Let us call the distillate nitric acid. If 
you have any nitric acid left after testing for its properties 
save it for the next experiment. 

Now remove the contents of the retort as follows: When 
the liquid in the retort has cooled somewhat, insert a funnel 
in the tubulature of the retort, and cautiously pour about 



FIRST YEAR CHEMISTRY 177 

100 c.c. of hot, distilled water directly into the midst of 
the liquid. If the contents of the retort are too hot, there 
is danger that the hot water will be converted into steam 
so rapidly that some of the hot acid mixture may be spurt- 
ed out from the retort, and if too cold, crystals may fix them- 
selves so firmly to the sides of the retort that they cannot 
be removed without danger of breaking the glass. When 
the water has all been added, shake until all crystals that 
may have been formed in the retort are dissolved and then 
empty the liquid contents of the retort into an evaporating 
dish and evaporate to crj/stallization. Examine the crys- 
tals closely and try to recognize them — by color, taste, form, 
solubility in water, and flame coloration — as the same sub- 
stance that you made in Experiment 71. If the resulting 
mixture in the evaporating dish is too strongly acid, or if 
the crystals are not satisfactory, free them from the strongly 
acid residual liquor and recrystallize them from hot water. 
What three simple substances are there in the crystals? 
Whence came the sulphur and the oxygen to form the po- 
tassium sulphate? In this reaction there were only two 
factors and only two products. Suppose we try to write 
the equation for the reaction between niter and sulphuric 
acid, using common names for those substances whose com- 
position we do not know. It will look as follows: 

} ( hydrogen \ J potassium 

j niter - + - sulphur V = J - d V + j sulphur 

( oxygen ) } ( oxygen 

Could the potassium that formed the potassium sulphate 
have come from the sulphuric acid? From which factor 
must it have come then? What, then, is one simple sub- 
stance in niter? Recall the action of sulphuric acid on 
iron sulphide and the action of sulphuric acid on sodium 
chloride and note the resemblance between those reactions 



178 FIRST YEAR CHEMISTRY 

and the one we are now studying. Was the hydrogen of 
the sulphuric acid evolved as a gas in the present experi- 
ment? What, then, became of the hydrogen that was in 
the sulphuric acid? Having traced some of the elements 
in our unknown substances, let us now rewrite the equa- 
tion as follows, letting X stand for the part that still re- 
mains unknown: 

(potassium | + j^rl = {Imogen | + \^T } 
' ' [ oxygen j L J ( oxygen j 

(niter) (sulphuric acid) (nitric acid) (potassium 

sulphate) 

Refer to the equation for the reaction between sodium 
chloride and sulphuric acid in Experiment 65 and note the 
similarity between that equation and the one last written. 
We are still in the dark as to what is united to the hydro- 
gen in nitric acid. We shall determine X in one of the fol- 
lowing experiments. Let us first see if we can actually find 
the hydrogen in nitric acid. 

Experiment 88. The action of magnesium on nitric acid. 
Have ready the test tube generator with short delivery 
tube used in Experiment 66 and illustrated in Fig. 64, gradu- 
ate, large beaker, small test tube, Bunsen burner, evapo- 
rating dish, tripod, and gauze; also nitric acid, and a piece 
of magnesium ribbon about 30 cm. long. 

Set up the test tube generator, fitted with a one-hole 
cork and a very short delivery tube leading to a beaker of 
water which serves the purpose of a pneumatic trough. 
Invert a small test tube full of water in the pneumatic 
trough, and let it stand in the water ready for use when 
the gas is to be caught. Put in the generator test tube 
about 5 c.c. of concentrated nitric acid from the bottle on 
the shelf, and add about 30 c.c. of water. Put in a piece 



FIRST YEAR CHEMISTRY 179 

of magnesium ribbon about 30 cm. long, at once insert the 
cork, and catch the gas evolved in a test tube over the 
trough. Touch the test tube of gas to a flame. What gas 
was evolved? Write the equation for evolving hydrogen 
from nitric acid by means of magnesium, representing the 
unknown part of nitric acid by X, as in the last equation 
in the preceding experiment. To find the other product 
of the reaction, take about 10 c.c. of the liquid in the test 
tube, and evaporate it just to dryness in an evaporating 
dish. From sulphuric acid you got sulphates, from hydro- 
chloric acid you got chlorides, and from carbonic acid you 
got carbonates. Nitric acid gives nitrates. The residue 
from the evaporation is called magnesium nitrate, and con- 
sists, of course, of magnesium and whatever was united 
with hydrogen in nitric acid. What one simple substance 
have you now proved to be in nitric acid? If you have 
saved any of the nitric acid that you made yourself, try 
some of it with magnesium ribbon to see if it duplicates 
the reaction of magnesium with the nitric acid from the 
bottle on the shelf. 

Experiment 89. The action of copper on nitric acid. Have 

ready the flask generator used in preparing hydrogen, long 
delivery tube, four fruit jars, pneumatic trough, glass stir- 
ring rod, deflagrating spoon, wooden toothpicks, beaker, 
and Bunsen burner; also some copper turnings, nitric acid, 
phosphorus, lime-water, and litmus paper. 

Generation of gas. — Put about 50 grams of copper turn- 
ings in the flask. Be sure that the thistle tube nearly 
touches the bottom of the flask. Add just enough water 
to seal the thistle tube. Then add successive small por- 
tions of concentrated nitric acid down the thistle tube till 
there is a good evolution of gas. The flask at first becomes 
filled with a reddish brown gas, but the gas evolved soon 



180 FIRST YEAR CHEMISTRY 

becomes colorless. When it is nearly colorless, catch four 
jars of the evolved gas over the pneumatic trough. Have 
the jars as free from water as possible. Reject the first 
jar of gas. From the second jar get the properties of the 
colorless gas. Is it hydrogen? Note that on exposure to 
the air it forms a brown gas. Get as many properties as 
you can of this brown gas. 

Composition of the colorless gas. — In the third jar, which 
should have little or no water in it, burn about half a 
gram of phosphorus, having the phosphorus well on fire be- 
fore plunging it into the jar. Keep the cover off the jar 
as short a time as possible. Seal the jar as quickly as pos- 
sible. Note the formation of white phosphorus oxide. 
Whence did the oxygen come to form this oxide? What 
is one simple substance, then, in the colorless gas evolved 
from the nitric acid? The proportion by volume of oxy- 
gen in the colorless gas may be ascertained by opening the 
jar under water. How much water runs in? Test the 
residual gas in the jar with a burning toothpick, but do not 
hold the burning toothpick long in the jar. What gas is 
indicated by the extinguishing of the burning toothpick? 
Try a second test for carbon dioxide by wetting a clean 
glass rod in clear lime-water and holding the rod in the 
midst of the gas left in the jar. If the gas is carbon dioxide, 
the rod will become covered with a white coating of calci- 
um carbonate. What would you now say about the com- 
position of the gas left in the jar? Is it carbon dioxide? 

If no white coating appears on the glass rod, it shows 
that you have a gas which puts out a burning splinter, 
but which is not carbon dioxide. Such a gas does exist, 
and it has figured incidentally in one of our early experi- 
ments; when we burned phosphorus in a jar of air in Ex- 
periment 12 we found that the jar was filled about one 
fifth full of water when it was opened in the pneumatic 









FIRST YEAR CHEMISTRY 181 

trough, i.e., that air is about one fifth oxygen. The other 
four fifths consist of nitrogen, a colorless gas that puts out 
a burning toothpick but does not turn lime-water milky; 
this gas was also mentioned in the note on air and its com- 
position at the end of Experiment 12. It must have been 
nitrogen, then, that you had left in the jar in this experi- 
ment after you had burned phosphorus in the colorless gas 
evolved from nitric acid by means of copper. Of what 
two simple substances was the colorless gas in which you 
burned phosphorus composed? What is the proportion of 
oxygen and of nitrogen in this gas? Let us call it colorless 
oxide of nitrogen ; sometimes it is called nitric oxide, and 
sometimes nitrogen monoxide. Whence must the nitrogen 
and the oxygen of nitric oxide have come? The water 
that was put into the flask to seal the thistle tube did not 
enter into the reaction between the copper and the nitric 
acid. What two other simple substances have you now 
proved to be in nitric acid besides the hydrogen discovered 
in the preceding experiment? Of what three simple sub- 
stances, then, is nitric acid composed? 

Composition of the reddish brown gas. — When the colorless 
oxide of nitrogen turned brown in air, what was happening 
to it chemically? What two other simple substances have 
you studied that formed two oxides each? Write the equa- 
tion for the spontaneous oxidation of nitric oxide in air. 
To distinguish the brown gas from the colorless one it is 
called nitrogen dioxide or nitrogen peroxide. 

Composition of nitric acid. — Finally generate a little more 
colorless oxide of nitrogen and catch a fruit jar from half 
to two thirds full of the gas. When you have removed the 
jar from the trough, loosen the cover a little, still holding 
the jar inverted, and let nearly all the water trickle out; 
about 20 or .30 c.c. of water should be left in the jar. As 
the water runs out, air, of course, passes in; this turns the 



182 FIRST YEAR CHEMISTRY 

gas in the jar brown; shake this brown gas well with the 
water left in the jar and then test it with litmus paper. 
The resulting liquid is a dilute solution of nitric acid. What 
further light does this throw on the composition of nitric 
acid? Write the equation for the addition of nitrogen di- 
oxide to water with the formation of nitric acid. What are 
the two component parts of nitric acid? Now that you 
have found out definitely the composition of nitric acid re- 
write the equation for niter and sulphuric acid. Of what 
three simple substances is niter composed? 

Treatment of the residue in the flask. — Decant the residual 
liquid from the flask into a beaker. The decantate should 
be clear, but it will probably be deeply colored. Set it 
away to crystallize. Get all the properties you can of 
crystallized copper nitrate. 

When you treated copper with nitric acid a colorless gas 
was evolved. What was it? A crystalline compound was 
obtained from the residual solution. What was it? There 
was also a third product whose presence we cannot show 
easily, and this product was water. Now write the equa- 
tion for copper and nitric acid, using as many portions of 
the acid as are necessary to give both nitric oxide and the 
nitrate of copper. Write the formula for nitric acid so as 
to show its component parts in full; also write the formula 
for copper nitrate to correspond thereto. 

What three compounds of nitrogen have you made in 
this experiment ? How many of them are gases ? 

Experiment 90. Reaction between potassium hydroxide 
and nitric acid. Have ready several beakers, graduate, 
glass stirring rod, Bunsen burner, tripod, gauze, horn-pan 
balance, set of smaller weights, and forceps; also some potas- 
sium hydroxide, nitric acid, litmus paper, and turmeric paper. 

Dissolve 10 grams of potassium hydrate in 30 c.c. of 



FIRST YEAR CHEMISTRY 183 

water. Dilute 10 e.c. of concentrated nitric acid with 30 c.c. 
of water. Neutralize the two solutions, evaporate to crys- 
tallization, and get the properties of the resulting compound, 
which is called potassium nitrate. Recognize it, by color, 
form, taste, solubility in water, and flame coloration, as the 
same substance as the niter started with in the investiga- 
tion of niter. Write the equation for the neutralization of 
potassium hydroxide and nitric acid. 

We have now decomposed potassium nitrate and we have 
made it by metathesis, thus completing our investigation. 
Of what three simple substances is nitric acid composed, 
and of w 7 hat three simple substances is niter composed? 

Theory of Chemistry. — Having now become well acquaint- 
ed with a large number of chemical substances, let us turn 
our attention to a brief discussion of the theories of chem- 
istry, and then continue our study of chemical substances, 
but in the light of present theoretical conceptions. 

Definition of Chemistry. — Thus far in the course we have 
been studying chemistry without trying to formulate a 
definition for it. We have made compound substances, 
we have broken down compound substances into simpler 
ones, we have studied the reaction between different sub- 
stances, we have noted the properties of many sub- 
stances, and we have studied the different kinds of chemic- 
al changes. All this may be condensed into the following 
statement: Chemistry deals with the composition of sub- 
stances, with changes of substances, and with phenomena at- 
tendant upon such changes. A consistent presentation of 
the subject should, of course, treat also of the laws govern- 
ing the different kinds of chemical changes, and the theories 
that have been proposed from time to time in explanation 
of the laws. It must be remembered that those laws which 
are based on mathematical data have been deduced from 



184 FIRST YEAR CHEMISTRY 

much experimental work, and that they are, to the best of 
our knowledge, unchanging facts; for that reason the laws 
of chemistry will be considered carefully. On the other 
hand, it must be remembered that the theories are not facts, 
but are simply attempted explanations to account for the 
existence and operation of the laws. It is natural, then, 
that the theories should change as the amount of knowl- 
edge concerning chemical changes increases. The older 
theories will be treated very briefly; they should not be 
omitted entirely, however, because they portray the trend 
of human thought in regard to the composition of matter, 
and they aid in a more intelligent comprehension of the 
modern theories. The theories held to-day by the majority 
of workers in natural science are considered the nearest 
approach, from the natural science point of view, to the 
true explanation of matter that mankind has yet reached, 
but it must be held constantly in mind that they may easily 
and rightfully be replaced to-morrow by other and more 
correct theories. Such is the law of progress till the goal of 
human endeavor is reached. 

Periods in the history of chemistry. — In studying the laws 
and theories of chemistry it is best to take them up chrono- 
logically. For our purpose the history of chemistry divides 
itself roughly into the following periods: 

I. Period of Alchemy. (Earliest times-1500). 

II. Medical Period. (1500-1650). 

III. Period of Robert Boyle. (1650-1700). 

IV. Phlogiston Period. (1700-1750). 
V. Pneumatic Period. (1750-1800). 

VI. Atomic Period. (1800-1900). 

VII. Modern Period. (1900-present time). 

The dates given above are only approximate, for in many 
cases the periods overlap each other a little. 



FIRST YEAR CHEMISTRY 185 

The Period of Alchemy extended from the earliest times 
to the end of the fifteenth century. The names of the 
three men most closely associated with the period of Al- 
chemy are Aristotle, Geber, and Valentine. The workers 
of this period are usually spoken of as the philosophers; 
their work did not consist of experiments planned to give 
an insight into the constitution of bodies; it consisted 
rather of speculations as to the nature of the world and 
the matter of which it consists. 

Aristotle was the most famous of the philosophers and 
he proposed that the world was made up of the four ele- 
ments, earth, water, air and fire. To Aristotle these words 
meant not what they mean to us, but the foundation sub- 
stances of which the earth is composed. Earth stood for that 
which is cold and dry; water for cold and wet; air, hot and 
wet ; fire, hot and dry. He soon saw that these four elements 
were not sufficient to explain all the phenomena; hence, he 
added a fifth, which he called cether, or the soul of matter. 

Geber was dissatisfied with the five elements of Aristotle; 
therefore, he added two more, mercury, which indicated 
fluidity and luster, and sulphur, which indicated hardness 
and combustibility. 

Valentine added still another element, salt, which indi- 
cated resistance to fire. 

The main pursuit of the alchemists was the transmuta- 
tion of metals, i.e., the changing of a base metal like lead 
to a noble metal like gold or silver. This pursuit led them 
to still another, the search for the Philosophers Stone, a 
substance of miraculous powers; not only was it to be 
the means by which the transmutations themselves were 
to be effected, but it was supposed to be the elixir of life 
and a cure for all disease. The following metals were known 
to the alchemists: gold, silver, tin, iron, copper, lead, ar- 
senic, antimony, bismuth, and zinc. 



186 FIRST YEAR CHEMISTRY 

The Medical Period began about 1500, and lasted about 
one hundred and fifty years. The names of four men are 
associated with this period, Paracelsus, Libavius, Van Hel- 
mont and Glauber. 

Paracelsus (1493-1541) joined the subjects of chemistry 
and medicine, with the result that new and better methods 
of preparing compounds were established and the proper- 
ties of chemicals were more carefully studied. The linking 
of the two subjects was the natural outgrowth of the 
alchemistic conception of the Philosopher's Stone; it should 
be borne in mind, however, that there is at the present 
day little or no connection between the two subjects; chem- 
istry deals with the composition of matter in general, medi- 
cine deals with the effect of a limited number of substances 
on the human system. 

Libavius (1540-1616) made sulphuric acid by burning 
sulphur and saltpeter, and showed that this was contained 
in the various crystalline vitriols. He also wrote the first 
clear and orderly arranged text-book on chemistry. 

John Baptist Van Helmont (1577-1644), a Dutch chem- 
ist, is noted for two things. First, he rejected the elements 
of Aristotle, Geber, and Valentine; unfortunately, he put 
nothing in their place, but he considered water as the pri- 
mary element. Secondly, he distinguished clearly between 
gases and vapors — calling aeriform substances which, when 
cooled, became liquids, vapors, and those which did not, 
gases; this distinction holds to-day. 

Johann Rudolph Glauber (1604-1666), a German chemist, 
is noted for his discovery of hydrochloric acid and sodium 
sulphate; the first he obtained by treating table salt with 
sulphuric acid; the second was obtained as a side product, 
and is called, in honor of his discovery, Glauber's Salt. 
He improved the methods for making nitric acid and glass. 
He also urged Germany to utilize its own natural resources 



FIRST YEAR CHEMISTRY 187 

for the production of new chemicals; in this he laid the 
foundations for modern technical and industrial chemistry. 

The Period of Robert Boyle began about 1650 and lasted 
about fifty years. As the name indicates, only one man, 
Boyle, was prominently identified with this period. 

Robert Boyle (1626-1691) was an English philosopher 
opposed to alchemy. His work helped to establish chem- 
istry as a separate science, independent of medicine, physics, 
and other subjects. First, he defined clearly the terms 
element, compound and mixture, and his definitions are 
nearly identical with those employed to-day; he noted the 
existence of about forty elements. Secondly, he estab- 
lished Qualitative Analysis on practically those principles 
that underlie that study to-day; he did not consider it nec- 
essary to isolate an element from a compound in order to 
prove the presence of that element; he found it more con- 
venient to show its presence by means of some reaction 
characteristic of that element; the test for certain metals 
by flame coloration is a good example. Thirdly, he stud- 
ied the effect of pressure on the volume of a gas and found 
that this may be expressed mathematically. Boyle's Law 
may be stated thus: The volume of a gas is inversely pro- 
portional to the pressure to which it is subjected. A later ex- 
periment will consider this law in detail. Lastly, he pro- 
proposed the Corpuscular Theory — that all matter exists in 
small particles, to which he gave the name corpuscles, and 
that compounds were made by the union of these particles. 
This theory of Boyle was a forerunner of the atomic theory. 

The Phlogiston Period lasted from about 1700 to about 
1750. Only one man, Stahl, was prominently connected 
with this period. 

George Ernst Stahl (1660-1734), a German physician and 
chemist, proposed that all combustible substances contain 
phlogiston; when they burn it goes out of them; when the 



188 FIRST YEAR CHEMISTRY 

oxides are reduced it comes back; carbon and sulphur were 
considered nearly pure phlogiston. The reasonableness of 
this theory is due to the fact that the use of the balance 
had not yet been introduced; hence, it was not recognized 
that when a substance burns there is an increase in weight. 
When this fact was pointed out later, the Phlogistonists 
contended that phlogiston had a negative weight; and 
when hydrogen was discovered later they considered this 
to be the long looked for phlogiston. The phlogiston theory, 
tho wrong, clung very tenaciously long after its falsity had 
been proved. Stahl's knowledge of the acids was more 
thoro than that of his predecessors. He recognized sul- 
phurous acid, made by burning sulphur, as different from 
sulphuric acid or oil of vitriol. 

The Pneumatic Period lasted from about 1750 to about 
1800. It was devoted largely to the study of gases and 
the composition of the atmosphere. But work on gases 
was not confined to this period by any means; it began 
with Van Helmont in the Medical Period and was contin- 
ued by Boyle. Even in our own day much important 
work has been done on aeriform matter, particularly in the 
liquefaction of gases. The men associated with the Pneu- 
matic Period are Cavendish, Black, Rutherford, Priestly, 
Bergman, Scheele, and Lavoisier. 

Henry Cavendish (1731-1810), an Englishman, is often 
called the Father of Pneumatic Chemistry. He discovered 
hydrogen in 1766 and called it " inflammable air"; later 
he discovered that a mixture of hydrogen and oxygen, when 
exploded, formed water. He saw the necessity of obtain- 
ing the exact specific gravities of gases in order to distin- 
guish one gas from another, and he himself determined 
many of these specific gravities. 

Joseph Black (1728-1799), a Scotchman, a physicist and 
a chemist, discovered carbon dioxide by heating marble, 



FIRST YEAR CHEMISTRY 189 

and called it " fixed air." He proved the difference be- 
tween carbon dioxide and air, and found that his " fixed 
air" could be taken up by " caustic alkalis" (our caustic hy- 
drates) making "mild alkalis" (our carbonates). He de- 
voted much time to experiments upon heat, discovering 
latent heat. He was probably the first man to raise bal- 
loons with hydrogen. 

Daniel Rutherford (1749-1819) was a Scotchman and a 
pupil of Black. He discovered nitrogen in 1772. 

Joseph Priestley (1733-1804) was an Englishman, first a 
clergyman and then a chemist. He discovered oxygen in 
1774 by heating red oxide of mercury. He failed to find the 
true explanation for combustion, but he reached the right 
conclusion in regard to the atmosphere, i.e., that air is "a 
mixture of two elastic fluids." He invented the pneumatic 
trough for collecting gases in practically the form used to- 
day. He also discovered nitrogen dioxide and carbon mon- 
oxide. 

Torbern Bergman (1735-1784) was a Swedish chemist 
and mineralogist. He developed and improved Boyle's 
methods of analysis, introducing many methods still used 
in Quantitative Analysis. He contributed much to the in- 
dustrial development of Sweden, working particularly on 
iron and steel. 

Carl Wilhelm Scheele (1742-1786) was a poor Swedish 
apothecary, but a chemist and a great experimenter. His 
life was filled with brilliant achievements, due probably to 
the fact that he made every move count, planning his ex- 
periments to arrive at definite results with the least ex- 
penditure of energy. Yenable in his History of Chemistry, 
says of Scheele: "No one, before nor since his day, has 
made so many important discoveries." 

He made many organic acids: lactic acid, from sour 
milk; tartaric acid, from cream of tartar; mucic acid, from 



190 



FIRST YEAR CHEMISTRY 



milk sugar; citric acid, from lemon juice; oxalic acid, from 
sawdust; malic acid, from apples; and gallic acid, from 
nut galls. He studied cyanogen, — a poisonous compound, 
composed of carbon and nitrogen, — and many of its deriva- 
tives, such as prussic acid and Prussian blue. Further- 
more, during his study of black oxide of manganese, — 
which was then called " black magnesia," — he found four 
substances, manganese, barium oxide, chlorine, and oxygen. 

Scheele's greatest work, perhaps, was the discovery, — in- 
dependently of Priestly, — of oxygen by heating red oxide 
of mercury. He also arrived at the right conclusion in re- 
gard to air, i.e., that it is " a mixture of two elastic fluids." 

Antoine Laurent Lavoisier (1743-1794), a Frenchman, was 
the central figure of the Pneumatic Period. His deep in- 
terest in physics led him to apply the balance to chemical 
problems, and this enabled him later to explain combustion 
aright. On account of his great service to chemistry he is 
sometimes called the Father of Chemistry. His most val- 
uable services were as an interpreter of the results of his 
own work and that of others. 

He overthrew the alchemistic belief that matter could 
be transmuted. This he did by boiling water for several 
days in a closed vessel, having first weighed the vessel and 
the water. Considerable earthy matter appeared in the 
water, but he proved that this was not " water" changed to 
"earth," by showing that the gain in weight of water cor- 
responded with the loss in weight of the vessel. 

He overthrew the phlogiston theory by his explanation of 
combustion. 

He formulated the Combustion Theory or Oxidation The- 
ory: (a) that substances burn only in pure air; (b) that 
this air is consumed in the combustion, and the increase in 
weight of the substance burned equals the decrease in weight 
of air; and (c) that the combustible body is, as a rule, con- 



FIRST YEAR CHEMISTRY 191 

verted into an acid by its combustion in pure air, but the 
metals are converted into metallic "calces," dull powders 
that are not acids. 

He proved the volumetric composition of the air by burn- 
ing phosphorus in a closed vessel and noting that four fifths 
of the air were not used in the burning. 

All this work of Lavoisier enabled him to formulate the 
Law of Conservation of Mass, or as it is sometimes called, 
the Law of Indestructibility of Matter: The sum of the 
weights of the products of a chemical change is exactly equal 
to the sum of the weights of the factors. This generalization 
is the first of four definite statements relating to the prin- 
ciples underlying chemical change; hence, it is often spoken 
of as the First Great Law of Chemistry. A later experiment 
will consider this law in detail. 

The Atomic Period lasted approximately from 1800 to 
1900. The division of matter into atoms was conceived in 
the first half of the century; the following years of the cen- 
tury were devoted to the study of the nature of the atoms. 
Many more men were associated with this period than with 
any preceding period, partly because the period was longer, 
and partly because the science of chemistry was taking 
definite shape and arousing the interest of students gener- 
ally. From the long list of workers we shall select the fol- 
lowing for further consideration: Richter, Berthollet, Proust, 
Dalton, Gay-Lussac, Avogadro, Mitscherlich, Dulong, Petit, 
Prout, Stas, Berzelius, Liebig, Hofmann, Woehler, Bunsen, 
Dumas, Davy, Faraday, Moissan, Meyer, and Mendeleeff. 

Jeremias Benjamin Richter (1762-1807), a German chem- 
ist, showed by experiment that the ratio of potassium hy- 
drate to sodium hydrate required to neutralize equal por- 
tion- of an acid was always the same, no matter what acid 
was used. This led him to believe that the composition of 
a substance is always the same no matter how the substance 



192 FIRST YEAR CHEMISTRY 

is made, — a belief that ripened later into one of the important 
statements of the period. 

Claude Louis Berthollet (1748-1822), a French chemist, 
did not agree with Richter, and contended that the compo- 
sition of a substance varied according to the varying amounts 
of the factors used. His work on lead oxide led him to 
think that metals form oxides with the addition of gradu- 
ally and indefinitely increasing amounts of oxygen. 

Joseph Louis Proust (1755-1826), another French chem- 
ist, proved by experiment that Berthollet's lead oxides 
were mixtures of the several oxides of lead, each of which 
was a distinct compound. He established the fact that 
lead and iron oxidize by jumps, and showed the difference 
between oxides and hydroxides. His controversy with 
Berthollet proved Richter to be right; and what Richter 
saw only indefinitely, Proust formulated as the Law of 
Definite Proportions by Weight: Every distinct chemical 
compound has a fixed and unalterable composition. This is 
generally called the Second Great Law of Chemistry, and it 
underlies that important branch of chemistry called Quali- 
tative Analysis. A later experiment will consider this law 
in detail. 

John Dalton (1766-1844), an English chemist, physicist, 
and mathematician, is the central figure of the Atomic 
Period. Tho sometimes weak in his experimentations he 
was good at drawing conclusions and formulating general 
laws. 

He formulated the Law of Expansion of Gases by Heat: 
The volume of a gas increases or decreases by one two hundred 
and seventy third of its volume as measured at 0° Centigrade 
for every degree rise or fall in temperature. Sometimes the 
law is stated as follows: The volume of a gas varies directly 
as the temperature on the u absolute scale." A later experi- 
ment will consider this law in detail. This law is some- 






FIRST YEAR CHEMISTRY 193 

times called the Law of Dalton and sometimes the Law of 
Charles, each man having formulated the law independent- 
ly of the other; tho it is of much importance, it is not con- 
sidered one of the great laws of chemistry, because it deals 
with a physical property of a gas rather than with its chem- 
ical composition. 

Dalton also discovered that water vapor exerts pressure, 
the same as any gas does but to a less extent; he called 
this vapor pressure vapor tension or the tension of aqueous 
vapor, and showed that it might be measured by a mercury 
column much the same as the pressure of the air is meas- 
ured by the hight of a mercury column. Vapor pressure 
must be taken into account in all experiments where the 
exact volume of a gas is to be determined. 

From his study of the two oxides of carbon and the five 
oxides of nitrogen Dalton formulated the Law of Multiple 
Proportions by Weight : When varying amounts of one sub- 
stance join a fixed quantity of some other, the varying amounts 
of the first bear to each other a simple numerical ratio, such as 
1:2, 1:3, 2:3, etc. This is considered the Third Great Law 
of Chemistry. A later experiment will consider this law in 
detail. 

Dalton formulated the Atomic Theory, the essence of 
which is that all matter is composed of small indivisible 
particles called atoms, which unite to form molecules. 
More in detail the theory may be stated thus: (a) Defini- 
tion of atoms : Every simple substance is made up of minute 
atoms, all alike, all of the same weight, and each one the small- 
est particle of that simple substance that can enter into com- 
bination with other atoms, (b) Definition of molecules : 
Every compound substance is made up of minute molecules, 
all alike, all of the same weight, and each a collection of atoms 
of different simple substances, grouped in simple and unalter- 
able numerical ratio and cJiemically combined, (c) A mole- 



194 FIRST YEAR CHEMISTRY 

cule is the smallest particle of matter that can exist by it- 
self. In a compound it is made up of atoms of different 
kinds. In a simple substance it is made up of atoms of the 
same kind. Often the molecule of a simple substance con- 
sists of only one atom. The subject matter of this para- 
graph, even if a little blind at the first reading, will become 
clearer and more meaningful when the student takes up 
the subject of the language of chemistry a little later. The 
atoms and molecules have never been seen, and probably 
never will be seen on account of their extremely small size. 
It has been calculated mathematically that if a drop of water 
were magnified to the size of the earth, the separate mole- 
cules of water in the drop would vary roughly in size from 
base-balls to foot-balls. 

Furthermore, Dalton thought he could find the relative 
weights of the atoms, i.e., the weight of one element that 
would unite chemically with unit weight of some other ele- 
ment taken as a standard. These relative weights he called 
combining numbers or atomic weights. He made a list of 
the combining numbers of several elements, referring them 
to hydrogen, the lightest known element. A later experi- 
ment will consider the subject of combining numbers in de- 
tail. 

Finally Dalton introduced a set of symbols to represent 
the atoms and their combination to form molecules. Some 
of the commoner ones and their combinations were repre- 
sented as follows: 

O represented an atom of oxygen. 

O represented an atom of hydrogen. 

• represented an atom of carbon. 

O O represented a molecule of water. 

O • represented a molecule of carbon monoxide. 

O • O represented a molecule of carbon dioxide. 



FIRST YEAR CHEMISTRY 195 

This clumsy system was the forerunner of our present, 
more convenient system. 

Joseph Louis Gay-Lussac (1778-1850), a French chemist 
and physicist, studied gases, — particularly the volumes of 
gases that enter into a chemical reaction. He formulated 
the Law of Definite Proportions by Volume : In any chemical 
change the relative volumes of the gaseous factors and products 
bear to each other a simple numerical ratio, such as 1:2, 1:3, 
2:3, etc.; incidentally, this ratio is always expressed by the 
number of molecules of the gaseous factors and products in the 
written equation. This law is called the Fourth Great Law of 
Chemistry. This law will be considered in qletail a little later. 

Amadeo Avogadro (1776-1856), an Italian physicist and 
chemist, suggested that "equal volumes of all substances, 
when in the state of gas, and under like conditions, contain 
the same number of molecules^ This statement cannot 
be proved by experiment, because the molecules cannot be 
seen or handled. Still Avogadro's Suggestion is probably 
correct, because it has enabled chemists to predict certain 
results from chemical experiments, and subsequent experi- 
ment has shown the predictions to be correct. The sug- 
gestion has proved most useful in determining the true 
atomic weights. 

Eilhard Mitscherlich (1794-1863), a German chemist, dis- 
covered that to a certain extent there existed a relationship 
between the crystalline forms of substances and the num- 
ber of atoms in the molecules of those substances. 

Pierre Louis Dulong (1785-1838), a French chemist and 
physicist, and Alexis Therese Petit (1791-1820), a French 
physicist, formulated what has since been called the Law 
of Dulong and Petit: The atoms of the different elements 
have the same capacity for heat. The application of this law 
will be considered in detail when we have determined some 
atomic weights. 



196 FIRST YEAR CHEMISTRY 

William Prout (1785-1850) 7 an English physician, thought 
that there was one fundamental substance, possibly hydro- 
gen, and that the different elements were condensations of 
this one substance. Hence, the combining numbers of all 
the elements should be multiples of that hydrogen, i.e., 
should be whole numbers. The proposition was very plaus- 
ible, but the atomic weights do not yet bear it out. 

Jean Servais Stas (1813-1891), a Belgian chemist, deter- 
mined very accurately the atomic weights of many ele- 
ments. His work is still considered authoritative, and his 
results indicate that Prout's hypothesis is probably incor- 
rect. 

Johann Jacob Berzelius (1779-1848), a Swedish chemist, 
determined many atomic weights with great care. He ana- 
lyzed minerals chemically, and discovered several rare ele- 
ments. He proposed an electrochemical theory, i.e., that 
every atom carried a certain quantity of electricity, partly 
positive and partly negative. He replaced Dalton's sym- 
bols by more convenient ones by proposing that the first 
letter of the Latin name of an element should represent one 
atom of that element. He gave the name allotropism to the 
existence of an element in different forms. 

Justus von Liebig (1803-1873), a German chemist, studied 
organic substances, i.e., those substances produced by va- 
rious forms of life. He established organic chemistry on a 
firm basis as a distinct branch of chemistry. Since his 
time it has grown so rapidly in size and importance that 
the most we can do in this course is to touch very briefly 
upon a dozen or so of the most important organic substances. 

August Wilhelm von Hofmann (1818-1892), a German 
chemist, devoted himself to an exhaustive study of organic 
chemistry, his work leading up to the coal-tar industry. 

Friederick Woehler (1800-1882), a German chemist, 
showed that organic and inorganic chemistry were not theo- 




Robert Bunsen 
(1811-1899) 



V\<;. W 




John Dalton 

(1766-1844) 



Fig. 69. Some noted chemists. 




Friederick Woehler 
(1800-1882) 






Antoine Laurent Lavoisier 
(1743-1794) 



Fig. 70. Some noted chemists. 




DlMITRI IVANOVITCH MENDELEEFF 

(1834-1907) 



Fig. 71 



FIRST YEAR CHEMISTRY 197 

retieally two distinct and separate subjects. He made sev- 
eral organic compounds from inorganic substances. 

Robert Bunsen (1S1 1-1899), a German chemist, invented 
the burner, the suction pump, and the electric battery that 
bear his name. He analyzed gases, and developed gas 
analysis as a separate branch of chemistry. He also intro- 
duced spectrum analysis, i.e., the detection of metals by 
their flame colorations. 

Jean Baptiste Andre Dumas (1800-1884), a French chem- 
ist, devised methods for determining the specific gravities of 
vapors and gases at high temperatures. 

Sir Humphrey Davy (1778-1829), an English chemist, 
studied the action of the electric current on water and on 
many other substances. He produced metallic sodium, po- 
tassium, barium, and calcium by electrolysis. He also stud- 
ied gases, particularly nitrous oxide. 

Michael Faraday (1791-1869), an English chemist and 
physicist, liquified chlorine and other gases. He also showed 
that when the same electric current passes thru separate solu- 
tions of metallic salts the amounts of the different metals de- 
posited are in the same ratio as the equivalent or combining 
weights of those metals. This is now known as Faraday's 
Law, and is much used in electrochemistry. 

Henri Moissan (1852-1907), a French chemist, devised the 
electric furnace, and made thereby many new substances, 
including artificial diamonds. 

Lothar Meyer (1830-1895), a German chemist, arranged 
the elements in a table according to their increasing atomic 
weights, and showed that similar properties recurred peri- 
odically thruout the list. This list is called the Periodic 
Table. 

Dimitri Ivanovitch Mendeleeff (1834-1907), a Russian 
chemist, formulated the Periodic Law: The properties of any 
dement arc periodic functions of its atomic weight. This law 



198 FIRST YEAR CHEMISTRY 

made possible the prediction of certain elements at the time 
unknown, and these were discovered in later chemical re- 
search. This periodic relationship will be considered more 
fully when we have become acquainted with a few more 
compounds. 

The language of chemistry. — In writing equations for the 
reactions that take place between chemical substances, we 
use symbols largely for the sake of saving time. In the 
equations that we have written so far, it was allowable to 
use contractions for the names of the simple substances, 
and we seldom represented the amounts used, because we 
were not studying the quantitative composition of the sub- 
stances. Now, however, since we have formed a definite 
conception of the atoms and of the molecules, it is desirable 
to make the equation show the "internal mechanism of the 
reaction," i.e., the change of place with respect to each 
other that the atoms undergo as they form new compounds. 

Symbols. — To represent an atom of a substance we use a 
symbol. Generally the first letter of the name of the element 
is used as the symbol; thus, H stands for the atom of hy- 
drogen, O for the atom of oxygen, S for the atom of sulphur, 
and C for the atom of carbon. If the names of two or more 
elements begin with the same letter, this initial letter is 
used as the symbol of one of the elements, and the symbol 
of the other element is made up of the initial letter and a 
second letter, which is characteristic of that name. Thus, 
C stands for carbon, Ca for calcium, CI for chlorine, Co for 
cobalt, and Cu for copper. Symbols consisting of one letter 
are always capitals. The first letter of a symbol consisting 
of two letters is always a capital and the second letter is 
always a small letter. Symbols are never followed by 
periods. 

In the case of a few of the metals which have been known 
from the earliest times the symbol is derived from the Latin 



FIRST YEAR CHEMISTRY 199 

name rather than from the English name. The symbol Cu 
for copper is an example, the Latin name for copper being 
cuprum. At present eighty-one elements are known. The 
table on page 200 contains the complete list of these elements 
with the symbol for each. The names of those which are 
rare are printed in italics, and it is not desirable to learn 
the symbols for these. The symbols of the more common 
elements will soon become familiar thru frequent use. The 
Latin name for each element whose symbol is derived from 
a Latin name is given in parentheses. 

Formulae. — To express a molecule of a compound we use a 
formula, i.e., a combination of the symbols of the elements 
contained in that compound. These symbols are written 
together without any connecting mark. Thus, CO stands 
for a molecule of carbon monoxide; HgO stands for a mole- 
cule of mercury oxide; ZnO stands for a molecule of zinc 
oxide; PbO represents a molecule of lead oxide; HC1 is 
the formula for a molecule of hydrochloric acid; and NaOH 
represents a molecule of sodium hydrate. In each of the 
examples just given there was in no case more than one 
atom of each simple substance in the molecule. Such a 
simple state of affairs must not always be expected. In 
many cases a molecule contains two or more atoms of the 
same kind; for instance, the molecule of nitric acid con- 
tains one atom of hydrogen, one of nitrogen, and three of 
oxygen. When there are several atoms of the same kind 
in the molecule, as in nitric acid just mentioned, subnumer- 
als are used; thus, HN0 3 is the formula for a molecule of 
nitric acid; CaC0 3 represents a molecule of calcium car- 
bonate; H 2 S0 4 represents a molecule of sulphuric acid; 
H 2 stands for a molecule of hydrogen gas, and 2 for a mole- 
cule of oxygen gas. The subfigure affects only the symbol 
immediately preceding it, and it always indicates the num- 
ber of atoms of that simple substance in a molecule of 



The Elements and their Symbols 



Aluminium .... 
Antimony (Stibium) 

Argon 

Arsenic ..... 

Barium 

Bismuth 

Boron 

Bromine 

Cadmium .... 
C cesium 



Calcium 

Carbon 

Cerium 

Chlorine 

Chromium 

Cobalt 

Columbium 

Copper (Cuprum) . . . 

Dysprosium 

Erbium 

Europium 

Fluorine 

Gadolinium 

Gallium 

Germanium 

Glucinum 

Gold (Aurum) .... 

Helium 

Hydrogen 

Indium 

Iodine 

Iridium 

Iron (Ferrum) .... 

Krypton 

Lanthanum 

Lead (Plumbum) . . . 

Lithium 

Lutecium 

Magnesium 

Manganese 

Mercury (Hydrargyrum) 



Al 

Sb 

A 

As 

Ba 

Bi 

B 

Br 

Cd 

Cs 

Ca 

C 

Ce 

CI 

Cr 

Co 

Cb 

Cu 

Dy 

Er 

Eu 

F 

Gd 

Ga 

Ge 

Gl 

Au 

He 

H 

In 

I 

Ir 

Fe 

Kr 

La 

Pb 

Li 

Lu 

Mg 

Mn 

Hg 



Molybdenum 

Neodymium 

Neon 

Nickel 

Nitrogen 

Osmium 

Oxygen 

Palladium 

Phosphorus 

Platinum 

Potassium (Kalium) . . 
Praseodymium . . . . 

Radium 

Rhodium 

Rubidium 

Ruthenium 

Samarium 

Scandium 

Selenium 

Silicon 

Silver (Argentum) . . 
Sodium (Natrium) . . 

Strontium 

Sulphur 

Tantalum 

Tellurium 

Terbium 

Thallium 

Thorium 

Thulium 

Tin (Stannum) . . . . 

Titanium 

Tungsten (Wolframium) 

Uranium 

Vanadium 

Xenon 

Ytterbium 

Yttrium 

Zinc 

Zirconium 



. Mo 
. Nd 
. Ne 
. Ni 
. N 
. Os 
. O 
. Pd 
. P 
. Pt 
. K 
. Pr 
. Ra 
. Rh 
. Rb 
. Ru 
. Sm 
. Sc 
. Se 
. Si 
• Ag 
. Na 
. Sr 
. S 
. Ta 
. Te 
. Tb 
. Tl 
. Th 
. Tm 
. Sn 
. Ti 
. W 
. U 
. V 
. Xe 
. Yb 
. Yt 
. Zn 
. Zr 



200 



FIRST YEAR CHEMISTRY 201 

the compound. Thus, the formula H 2 S0 4 indicates that the 
molecule of sulphuric acid contains two atom of hydrogen, 
one atom of sulphur, and four atoms of oxygen. The 
proper number of atoms of each simple substance in a mole- 
cule of a compound must always be determined by a careful 
quantitative examination of the substance. Manifestly, we 
have not the time for such elaborate analyses. It would 
seem easier for us simply to commit the correct formulae 
to memory after they had been determined by analysis; 
even this laborious process may be much shortened b} r cer- 
tain helps to the writing of equations that are considered in 
the next few pages. 

Coefficients. — To express tiro or more rdolecules, coeffi- 
cients are used; thus, 5HN0 3 means five molecules of nitric 
acid; 2CaC0 3 means two molecules of marble; 6H 2 S0 4 
means six molecules of sulphuric acid; 2H 2 means two 
molecules of hydrogen gas; and 40 2 means four molecules of 
oxygen gas. Occasionally it is desirable to represent sev- 
eral separate atoms of an element not necessarily united to 
another substance nor to each other. Thus, 2C stands for 
two atoms of carbon, 4S stands for four atoms of sulphur, 
30 stands for three atoms of oxygen, and 5H stands for 
five atoms of hydrogen. 

It is important to remember that the subnumerals always 
show the number of atoms of a simple substance within a 
molecule; these atoms are always chemically united. A co- 
efficient, on the other hand, indicates the number of complete 
molecxdes with no chemical union between the separate mole- 
cules; the coefficient may occasionally represent a number 
of isolated atoms of a simple substance, but this use is of 
rare occurence. 

Writing equations in the new form, using symbols and 
formula', is not a difficult matter. The principles under- 
lying our old method of writing equations hold here, namely, 



202 FIRST YEAR CHEMISTRY 

the principle (1) of determining by actual experiment 
what factors and products actually take part in the reac- 
tion, (2) of showing the simple substances in every com- 
pound, (3) of putting factors to the left and products to the 
right of the equality sign, and (4) of making the equation 
balance. The main difference is that now we shall show the 
exact quantitative composition of every substance that en- 
ters into the reaction. The following examples may serve 
to illustrate the new method. 

The reaction between magnesium and sulphuric acid yield- 
ing hydrogen and magnesium sulphate is represented as 
follows: 

Mg + H 2 S0 4 = H 2 + MgS0 4 . 

The reaction between zinc oxide and sulphuric acid yield- 
ing zinc sulphate and water is represented as follows: 

ZnO + H 2 S0 4 = ZnS0 4 + H 2 0. 

The neutralization of potassium hydrate with hydrochloric 
acid with the production of potassium chloride and water is 
shown as follows: 

KOH + HC1 = KC1 + H 2 0. 

The neutralization of sodium hydrate with sulphuric acid 
with the production of sodium sulphate and water is written 
as follows: 

2NaOH + H 2 S0 4 = Na 2 S0 4 + 2H 2 0. 

The oxidation of copper in air is expressed by the equation: 

2Cu + 2 = 2CuO. 

Helps in writing equations. — For the easy writing of equa- 
tions, the following points are exceedingly helpful: 

(1) Some knowledge of Valence with the Key to the Va- 
lence of the common elements. 



FIRST YEAR CHEMISTRY 203 

(2) The list of the formulae of the eight common acids 
and the method of building up the formulae of their corre- 
sponding salt?. 

(3) The list of the elements that have two atoms to the 
molecule. 

These helps will be taken up in detail in the above order, 
and then the safest and best way to write an equation will 
be explained and illustrated. 

Valence ?'s the power an atom of an elementary substance 
has of holding other atoms. It is sometimes called Valency 
or Quantivalence. Valence is usually expressed numerically, 
but it is not customary to indicate the amount of the va- 
lence by a numeral or any other distinguishing mark. The 
valence of an element is always determined by an examina- 
tion of the formulae of some compounds containing that ele- 
ment. The formulae were not originally derived from the 
valence but from a consideration of the molecular weight 
of the substance in connection with its percentage composi- 
tion. If we examine the formulae of some of the compounds 
of hydrogen we find that chlorine combines with one atom 
of hydrogen, oxygen with two, nitrogen with three, and 
carbon with four, as shown in the following list: 

Hydrochloric Acid HC1 

Water H 2 

Ammonia H 3 N 

Marsh Gas H 4 C 

If we take hydrogen as our unit of power we might say that 
chlorine has a valence of 1, oxgyen a valence of 2, nitrogen 
a valence of 3, and carbon a valence of 4. But not many 
elements unite with hydrogen. To ascertain the valence of 
those elements which do not unite with hydrogen we must 
turn to some other class of compounds, such as the oxides. 
Careful quantitative analysis of the different oxides shows 



204 FIRST YEAR CHEMISTRY 

that the molecules of some of the more common ones are 
expressed by the following formulae: 



MgO 


PbO 


Na 2 


co 2 


As 2 3 


Pb 3 4 


ZnO 


FeO 


H 2 


so 2 


Sb 2 3 


A1 2 3 


CuO 


MnO 


K 2 


Si0 2 


p 2 5 


Fe 3 4 


HgO 


BaO 


Ag 2 


Mn0 2 


N 2 5 


Fe 2 3 


CaO 


NO 


N 2 


N0 2 


S0 3 


Bi 2 3 



If we examine this list with respect to the oxygen, we 
find that a large number of the elements unite with oxygen, 
atom to atom, i.e., the power of such metals as Mg, Zn and 
Cu to hold other atoms is the same as that of oxygen. In 
sodium oxide, in water, and in a few other oxides, each 
oxygen atom can hold two atoms of the other element, i.e., 
oxygen has twice as great a valence as sodium, hydrogen, 
etc. No substance is known more than two of whose atoms 
will unite with one atom of oxygen. As the hydrogen atom 
has the lowest atom-holding power known, it is often taken 
as a unit of measure and the valence of hydrogen is said to 
be 1. Other elements that like hydrogen unite with oxy- 
gen two atoms to one, are said to have a valence of 1, or 
are called univalent elements. Oxygen has a valence of 2, 
since one atom of it will hold two atoms of hydrogen, and 
all elements that unite with oxygen atom for atom have a 
valence of 2, or are called bivalent. 

In substances such as carbon dioxide and the first oxide 
of sulphur we find substances whose atom will hold two 
atoms of oxygen; carbon and sulphur under such circum- 
stances have a valence of 4, or are tetravalent. The oxides 
of arsenic and antimony show that these elements have a 
valence of 3, or are trivalent. Phosphorus and nitrogen in 
some compounds each have a valence of 5, or are pentava- 
lent. Sulphur in the second oxide has a valence of 6, or is 
sexivalent. 



FIRST YEAR CHEMISTRY 205 

The more common valences are 1, 2, and 3; higher valences 
occur only occasionally and they may be neglected, as it is 
easier in such cases to build up the formulae of salts from 
the formula? of the acids themselves, as will be shown shortly. 
The valence of an element is generally constant, particularly 
in the case of metals. In a few cases the valence of an ele- 
ment varies with the conditions of the experiment, as in 
the case of sulphurous oxide and sulphuric oxide, where 
the valences are 4 and 6 respectively. Elements with 
varying valences are usually acid-forming substances and 
the confusion caused by this variation is obviated by means 
of the help referred to in the first sentence of this paragraph, 
e.g., building up the formulae of the salts from those of the 
acids. Elements of the same valence combine with or re- 
place each other atom for atom. Thus, one atom of potas- 
sium replaces one atom of hydrogen in hydrochloric acid, 
and one atom of oxygen combines with one atom of copper. 
Elements of different valence form compounds in which 
the number of valences belonging to one element exactly 
equals the number of valences belonging to the other ele- 
ment; the compound is then said to be " saturated" or to 
have no free valence. It goes without saying that the mole- 
cule of a compound has no valence, because all the atom- 
holding power of each element in the molecule is already 
satisfied in the complete molecule. 

Key to the valence of the elements. — As stated above, the 
valence of an element is usually constant. For convenience 
in determining the correct formulae of chemical compounds, 
the following list may be found of use: 

Univalent: H, Na, K, Li, (NH 4 ), F, CI, Br, I, Ag. 

Trivalent: As, Sb, Bi, Al, Au ; and sometimes Fe. 

Bivalent : Almost all the rest of the common metals ; 
also ; also S in sulphides. 

High valences may be disregarded, as it is easier to learn 



206 FIRST YEAR CHEMISTRY 

by memory the formulae of the few compounds wherein 
higher valences occur, or, if possible, to build up such for- 
mulae as previously suggested. 

Structural formulae. — In an ordinary equation the va- 
lences of the elements are never shown, except in so far as 
the formulae themselves indicate the atom-holding power 
of each element. In certain cases it is convenient to illus- 
trate the valences by bonds or hands thus: 

Na— — Zn— — As= =C = 

but in such cases the atoms are always grouped in such a 
way that all the bonds of an element are satisfied or joined 
to the bonds of another element. 

A structural formula is a graphical representation of a mole- 
cule, showing each element in the molecule tied by its bonds 
to the other elements in such a way that they form a unit with 
all the bonds satisfied. Sometimes it is called a graphical 
formula. Some simple illustrations are: 

H— 0— H H— H Zn = Na— O— H Na— CI 

These stand for the molecules of water, hydrogen gas, zinc 
oxide, sodium hydrate, and sodium chloride respectively. 
The one line between the two hydrogen atoms in the struc- 
tural formula for hydrogen gas indicates that the unit va- 
lences of the two separate atoms have united to form one 
bond. In like manner in the structural formula for sodium 
hydrate the unit valence of the sodium atom has united 
with one of the valences of the oxygen atom to form a 
single bond, and the unit valence of the hydrogen atom has 
united with the other valence of the oxygen atom to form a 
single bond. Complex molecules may be pictured in the 
same way, but since there is some doubt as to which possible 
combination of atoms expresses the true internal structure 
of the molecule it is customary to leave the acid radicals 



FIRST YEAR CHEMISTRY 



207 



as groups and show the hydrogen or metallic part structur- 
ally, thus: 



H^(S0 4 ) 
Sulphuric acid 

H-(N0 3 ) 
Nitric acid 



Na- 
Na- 



(S0 4 ) 



Sodium sulphate 

K— (N0 3 ) 

Potassium nitrate 



Mg=(S0 4 ) 

Magnesium sulphate 



Ca 



(N0 3 ) 

(N0 3 ) 
Calcium nitrate 









H— (POJ 

Phosphoric acid 



Na\ 

Na— (P0 4 ) 
Na^ 

Sodium phosphate 



Ca Ca Ca 

w / \ // 
(po 4 ) (po 4 ) 

Calcium phosphate 



Structural formulae are a help principally in determining 
the correct formulae of molecules containing a large number 
of atoms or atoms of high valence, tho they are often used 
in picturing the structural formulae of the simpler molecules. 
How to determine the formula of a salt. — The easiest way 
to determine the correct formula for a salt is not to build 
it up from the simple elements with the aid of the valence 
of each, because it is sometimes possible to make a large 
number of combinations that will satisfy the valences, 
whereas a careful consideration of many compounds shows 
that each compound has one definite constitution. It is 
much better to learn the formulae of the eight common 
acids and then to replace the hydrogen by the proper metal, 
bearing in mind the valence of the metal. 
The eight common acids are : 

H 2 S0 4 sulphuric acid. HC1 hydrochloric acid. 

H 2 S0 3 Sulphurous acid. HN0 3 nitric acid. 

H 2 S hydrogen sulphide. H3PO4 phosphoric acid. 

HCIO3 chloric acid. H2CO3 carbonic acid. 



208 



FIRST YEAR CHEMISTRY 



In sulphuric acid each of the two hydrogen atoms is uni- 
valent and must be united thru its bond to the rest of the 
molecule: The (S0 4 ) group does not exist alone, but is called 
the sulphuric acid radical, and has a valence of 2, since it 
holds two hydrogen atoms; it can, therefore, unite with 
one atom of a bivalent metal or with two atoms of a uni- 
valent metal. In like manner we speak of (N0 3 ) as the 
nitric acid radical with its valence of 1 and (CO3) as the car- 
bonic acid radical with its valence of 2. The following list 
shows by means of a few examples how the principle can be 
applied to all metals: 



H 2 S0 4 



H 2 S0 3 H 2 S HC1 HNO3 H 2 C0 3 H 3 P0 4 



K 2 S0 4 


Na 2 S0 3 CuS 


AgCl 


AgN0 3 


CaS0 4 


K 2 S0 3 Ag 2 S 


BaCl 2 


Cu(N0 3 ) 2 


A1 2 (S0 4 ) 3 


CaS0 3 As 2 S 3 


BiCl 3 


Bi(N0 3 j 3 



Na 2 C0 3 Na 3 P0 4 
MgC0 3 K 3 P0 4 

A1 2 (C0 3 ) 3 Ca 3 (P0 4 ) 2 



In this way by learning the valence of fifteen elements 
and the formulae of eight acids it is possible to build up the 
proper formulas of all the common salts instead of commit- 
ing them all to memory. A short time ago it was stated 
that the formulas were not originally derived from the va- 
lence, but the valence from the formulas. It is allowable for 
us now, however, to use the valence in writing the formulas, 
because we have not the time to do all the experiments that 
determine the formulas, nor to memorize the formulas as 
determined by others. 

How to determine the formula of a hydrate. — Since al- 
most all metals form hydrates, it is desirable to be able to 
decide quickly upon the correct formula for the hydrate of 
the metal in question. All hydrates are considered as be- 
ing derived from water by the replacement of half the hy- 
drogen in the water by the metal. If the formula for water 
be written HOH, it is easy to see that the univalent medals 



FIRST YEAR CHEMISTRY 209 

will form the hydrates NaOH, KOH, and NH 4 OH, when 
only one hydrogen atom of the water molecule is replaced by 
the metal. The (OH) group is called hydroxyl, and its va- 
lence is, of course, 1. A univalent metal will unite with 
one hydroxyl, a bivalent metal will unite with two hy- 
droxyls, and a trivalent metal with three hydroxyls to form 
hydrates, e.g., KOH, Ca(OH) 2 , Mg(OH) 2 , and Al(OH) 3 ; 
the last three are sometimes written Ca0 2 H 2 , Mg0 2 H 2 , and 
AIO3H3, but the first set of formulae is preferable. 

The number of atoms in the molecule of an elementary 
substance. — The application of Avogadro's Suggestion to the 
volumes of gases shows that the three common gases, oxy- 
gen, hydrogen, and nitrogen, and the four halogens, fluo- 
rine, chlorine, bromine, and iodine, all have two atoms to 
the molecule. The proper formulae for these seven elements 
are, therefore: 

2 H 2 N 2 F 2 Cl 2 Br 2 and I 2 

and these should be used wherever the element occurs alone 
in an equation. All other common substances have only 
one atom to the molecule, except phosphorus, which has 
four atoms. In such cases 

Na K Cu Fe C S and P 4 

represent the molecules of these substances correctly. It is 
important not to confuse valence with the number of atoms 
in the molecule; the former is never shown in an equation 
by any distinguishing mark; the latter is always indicated 
by a subfigure placed after the symbol of the element. 

Proof that the molecule of hydrogen contains two atoms. — Experi- 
ment shows thai one volume of hydrogen unites with one volume of 
chlorine to form two volumes of hydrochloric acid gas. Let us as- 
sume that the volume of hydrogen contains 1000 molecules. Then, 
according to Avogadro's Suggestion, the equal volume of chlorine con- 
tains 1000 molecules of chlorine and the two volumes of the product 



210 FIRST YEAR CHEMISTRY 

contain 2000 molecules of hydrochloric acid gas. Since every mole- 
cule of hydrochloric acid gas contains at least one atom of hydrogen 
and at least one atom of chlorine, the 2000 molecules of the product 
must contain in all 2000 atoms of hydrogen and 2000 atoms of chlorine. 
Therefore, each molecule of hydrogen gas must be divisible into two 
atoms of hydrogen, since the 1000 molecules of hydrogen furnished 2000 
atoms of hydrogen. For a similar reason each molecule of chlorine 
must contain two atoms of chlorine. It may be proved by similar 
reasoning that each of the other members of the above list has two 
atoms to the molecule. 

How to write an equation. — There are four steps in the 
writing of an equation. 

(1) Determine by actual experiment , or by reference 
to the record of the experiment in the laboratory note- 
book, the number of factors and products entering into the 
reaction; determine also the chemical name of each factor 
and product. 

(2) Determine the correct formula for each factor and 
product, using the helps on valence, the number of atoms 
to the molecule in simple substances, and the building up 
of formulae of salts. 

(3) Put these correct formulge in the form of an equa- 
tion, the factors to the left and the products to the right of 
the equality sign. 

(4) Balance the equation by taking the proper number 
of molecules of each factor or product or both to give the 
same number of atoms of each element on both sides of the 
equation. Balancing should always be done by changing the 
number of molecules, if necessary, and never by changing the 
number of atoms in the molecular formulce, for the formula? 
have already been decided upon as right. 

In each case an equation should represent an experiment 
actually performed, as it is generally worse than useless to 
write an equation by simply making different combinations 
of atoms. 






FIRST YEAR CHEMISTRY 211 

Commit to memory the three helps that are printed in 
heavy faced type on the preceding pages, namely: (1) The 
key to the valences of the elements, (2) the list of eight com- 
mon acids, and (3) the list of seven elements that have two 
atoms to the molecule. 

Writing equations illustrated. — Let us now take the first 
equation given on Page 202, and run thru the steps neces- 
sary for writing this equation. The question w 7 ould proba- 
bly read as follows: Write the equation for the action of sul- 
phuric acid on magnesium. The four steps are: (1) The 
question states the factors. The laboratory notebook rec- 
ord of the experiment tells us that hydrogen was evolved, 
and that magnesium sulphate was left in solution. (2) The 
list of eight common acids tells us that H 2 S0 4 is the formula 
for sulphuric acid. Since Mg is a bivalent element, the for- 
mula for magnesium sulphate must be MgS0 4 . The for- 
mula for hydrogen gas must be written H 2 , and that for 
magnesium must be written Mg, because the former is in 
the list of elements having two atoms to the molecule and 
the latter is not in this list. (3) Putting the factors and 
products in the form of an equation we get: 

Mg + H 2 S0 4 = MgS0 4 + H 2 . 

(4) Balance the equation by counting the number of Mg 
atoms on the left and those on the right of the equality sign. 
Since there is one atom of Mg on each side, the equation 
balances with respect to the magnesium. Next count the 
number of hydrogen atoms on each side, then the sulphur 
atoms, and finally the oxygen atoms. It will be seen that 
the equation balances, because we have just as many atoms 
of each element on one side as on the other. 

Let us take another case, — the fourth one given on Page 
202. When we have finished the third step in the writing 



212 FIRST YEAR CHEMISTRY 

of the equation for the neutralization of sodium hydrate in 
sulphuric acid, we get this: 

NaOH + H 2 S0 4 = Na 2 S0 4 + H 2 0. 

When we balance this equation we see that the smallest 
number of sodium atoms we can have on the right hand side 
of the equality sign is two. That necessitates taking two 
molecules of NaOH, and we indicate this by putting a co- 
efficient 2 in front of the NaOH. We now have hydrogen 
atoms enough on the left hand side to form two molecules 
of water and we show this by putting a coefficient 2 in front 
of the H 2 0. The equation then looks like this: 

2NaOH + H 2 S0 4 = Na 2 S0 4 + 2H 2 0. 

By counting the number of atoms of each element on both 
sides of the equality sign we find that the equation balances. 

List of equations for Experiments 1 to 90. — Let us now re- 
view briefly the experiments already performed and make 
a complete list of the equations figuring therein, using sym- 
bols and formulae. For those experiments which deal with 
the properties of substances it is impossible, of course, to 
write equations. The numbers in the list given below in- 
dicate the number of the experiment. Put the complete 
list in the laboratory notebook. 

Write the following equations: 

6. Write the equation for heating copper in air. Here 
it is allowable to neglect the nitrogen of the air and to write 
the equation as if copper and oxygen were the only ele- 
ments present. Remember that oxygen is one of the list 
of seven elementary substances that always have two atoms 
to the molecule when the elements exist alone. 

8. Heating zinc in air. 
10. Burning magnesium in air. Note that the three 






FIRST YEAR CHEMISTRY 213 

metals already used are all bivalent and have but one atom 
to the molecule when in the elementary state. 

12. Burning phosphorus in air. Note that phosphorus 
is the one exception to the rule that all elementary sub- 
stances except seven have but one atom to the molecule. 
The molecule of phosphorus is written P 4 , because it has 
four atoms to the molecule. Furthermore, phosphorus is 
pentavalent; the molecule for phosphorus oxide is written 
P 2 5 ; this formula should be memorized. 

14. Heating iron in air. Iron may be considered as bi- 
valent, unless something is said to the contrary. The for- 
mula for the simplest iron oxide is, therefore, FeO. It should 
be noted, however, that when iron "rusts" or oxidizes spon- 
taneously in air, the red-brown oxide formed has the for- 
mula Fe 2 03: in this compound iron is trivalent. The formula 
for the oxide of iron obtained by heating iron in air is Fe 3 4 ; 
write the equation for heating iron in air; also write the 
equation for the rusting of iron. 

17. Heating red oxide of mercury in air. 

18. Preparation of oxygen from potassium chlorate by 
means of heat. Here refer to the second help for writing 
equations — the list of common acids — and note that the 
formula for chloric acid is HC10 3 . The key to the valence 
of the elements tells us that potassium is univalent; the for- 
mula for potassium chlorate must be KC10 3 . 

20. Burning zinc in oxygen. This equation is, of course, 
the duplicate of the equation in 8. 

21. Burning magnesium in oxygen. 

22. Burning phosphorus in oxygen. 

23. Decomposing water by electricity. Remember here 
that the two gases produced are obtained in the molecular 

te. 

24. Producing water from hydrogen and oxygen by 
means of the electric spark. 



214 FIRST YEAR CHEMISTRY 

25. Preparation of hydrogen from steam by means of 
magnesium. The formula for steam is the same as that 
for water. 

28. Burning sulphur in air or oxygen. The first oxide of 
sulphur is sulphur dioxide, S0 2 ; there is no sulphur monoxide. 

29. Producing sulphurous acid by adding water to sul- 
phur dioxide. 

30. Producing the second oxide of sulphur by passing 
oxygen and the first oxide of sulphur over hot platinum 
sponge; the second oxide of sulphur is sulphur trioxide, S0 3 . 

31. Producing sulphuric acid by adding water to sulphur 
trioxide. 

32(a). Preparation of hydrogen from sulphuric acid by 
means of zinc. The formula for zinc sulphate in solution 
may be written ZnS0 4 . 

32(b). Efflorescence of crystallized zinc sulphate. Care- 
ful analysis shows that for every molecule of zinc sulphate 
in the crystallized salt there are seven molecules of water of 
crystallization. The formula for the crystallized salt is 
written ZnS04.7H 2 0. The period in the middle of the for- 
mula indicates chemical union between the water of crystal- 
lization and the rest of the salt. Do not try to remember 
the number of molecules of water of crystallization in crys- 
tallized salts. 

32(c). Burning hydrogen in air. 

33(a). The action of sulphuric acid on iron. 

33(b). Efflorescing of crystallized iron sulphate, 
FeS0 4 .7H 2 0. 

34(a). The action of sulphuric acid on magnesium. 

34(b). Efflorescing of crystallized magnesium sulphate, 
MgS0 4 .7H 2 0. 

35(a). The action of sulphuric acid on copper. Ref- 
erence to the record of this experiment will remind you that 
the first oxide of sulphur instead of hydrogen was evolved, 



FIRST YEAR CHEMISTRY 215 

that copper sulphate was left in solution, and that there was 
a third product — water. 

35(b). Efflorescing of crystallized copper sulphate, 
CUSO4.0H2O. The number of molecules of water of crys- 
tallization varies with different salts, tho the three crystal- 
lized salts studied above each had seven molecules. 

36. Reaction between zinc oxide and sulphuric acid. 

37. Reaction between copper and sulphur with the pro- 
duction of copper sulphide. Sulphur in sulphides is always 
bivalent, as may be seen from the formula for hydrogen 
sulphide in the list of acids. 

38. Reaction between mercury and sulphur. 

39. Reaction between zinc and sulphur. 

40. Reaction between iron and sulphur. 

41(a). Preparation of hydrogen sulphide by passing hy- 
drogen thru sulphur vapors. Do not represent the zinc and 
the sulphuric acid in the equation. Start, instead, with the 
hydrogen in the gaseous state. 

41(b). Decomposition of hydrogen sulphide by heat- 
ing the gas in the capillary tube. 

42(a). Reaction of sulphuric acid and iron sulphide. 

42(b). Burning of hydrogen sulphide in air, assuming 
that the hydrogen and the sulphur both oxidize completely. 

43. Precipitation of copper sulphide by passing hydro- 
gen sulphide into copper sulphate solution. 

44. Reaction of sulphuric acid with zinc sulphide. 
46(a). Burning carbon in air or oxygen with the forma- 
tion of carbon dioxide, C0 2 . 

46(b). Preparation of carbonic acid by dissolving car- 
bon dioxide in water. 

46(c). Decomposition of carbonic acid into its compo- 
nent parts by means of heat. 

47. Complete reduction of carbon dioxide by means of 
magnesium. 



216 FIRST YEAR CHEMISTRY 

48(a). Reducing carbon dioxide to carbon monoxide, 
CO, by means of zinc. 

48(b). Burning carbon monoxide in air with the pro- 
duction of carbon dioxide. 

48(c). Reduction of carbon dioxide to carbon monox- 
die by means of carbon. 

50. Burning hydrogen in chlorine. 

51. Action of hydrochloric acid on zinc. 

52. Action of hydrochloric acid on iron. 

53. Action of hydrochloric acid on magnesium. 

55. Heating sodium in contact with air. Since low heat 
produces sodium monoxide, Na 2 and high heat produces 
sodium dioxide, Na 2 2 , it will be necessary to write two sepa- 
rate equations in this case. 

56(a). Preparation of sodium hydrate by adding water 
to sodium monoxide. 

56(b). Action of water on sodium dioxide. 

56(c). Burning iron in oxygen. 

57. Action of sodium on water. First write the twin 
equations that represent this change. Then show the 
whole change in one equation. 

58. Action of sodium with chlorine. 

59(a). Neutralization of sodium hydrate with sulphu- 
ric acid. 

59(b). Efflorescence of crystallized sodium sulphate, 
Na 2 SO 4 .10H 2 O. 

60. Neutralization of sodium hydrate with hydrochloric 
acid. 

61(a). Preparation of sodium carbonate from sodium 
hydrate and carbonic acid. 

61(b). Efflorescence of crystallized sodium carbonate, 
Na 2 CO 3 .10H 2 O. 

62. Preparation of sodium carbonate from sodium hy- 
drate and carbon dioxide. 



FIRST YEAR CHEMISTRY 217 

63. Reaction between sodium carbonate and sulphuric 
acid. 

64. Reaction between sodium carbonate and hydrochlo- 
ric acid. 

65. Reaction between sodium chloride and sulphuric acid. 

66(a). Preparation of sodium amalgam. Tho this sub- 
stance has not a definite composition it is allowable to write 
its formula NaHg in this experiment. To be strictly cor- 
rect one should write it Na x Hg y . 

66(b). Reaction between sodium amalgam and water. 

68. Oxidizing potassium. Only one oxide is formed, 
potassium monoxide, K 2 0. 

69. Preparation of potassium hydrate by adding water 
to potassium oxide. 

70. Action of potassium on water. First write the twin 
equations that represent this change. Then show the whole 
change in one equation. 

71. Neutralization of potassium hydrate with sulphuric 
acid. 

72. Neutralization of potassium hydrate with hydrochlo- 
ric acid. 

73. Preparation of potassium carbonate from potassium 
hydrate and carbonic acid. 

74. Production of potassium carbonate by passing car- 
bon dioxide over hot potassium. First write the twin equa- 
tions that represent this change. Then show the whole 
change in one equation. 

76. Heating calcium in contact with air. 

77. Preparation of calcium hydrate by slaking quick- 
lime. The formula for calcium hydrate may be written 
either Ca0 2 H 2 or Ca(OH) 2 , but the second is preferable. 

7s. Action of calcium on water. First write the twin 
equations that represent this change. Then show the whole 
change in one equation. 



218 FIRST YEAR CHEMISTRY 

79. Reaction between calcium hydrate and hydrochloric 
acid. 

80(a). Precipitation of calcium sulphate from calcium 
chloride solution by means of sulphuric acid. 

80(b). Driving water of crystallization out of gypsum, 
CaS0 4 .2H 2 0, by means of heat. 

80(c). The setting up of plaster of Paris when water is 
added to it. CaS04 may be regarded as the formula for plas- 
ter of Paris, and the setting up as the formation of minute 
crystals of gypsum, CaS0 4 .2H 2 0. 

80(d). Preparation of calcium sulphate by adding sul- 
phuric acid to. calcium hydrate. 

81(a). Precipitation of calcium carbonate from calcium 
hydrate solution by means of carbonic acid. 

81(b). Decomposition of carbonic acid by means of heat. 

82. Precipitation of calcium carbonate by passing car- 
bon dioxide into calcium hydrate solution. 

83(a). Decomposition of calcium carbonate by means of 
heat. 

83(b). Passing carbon dioxide thru lime water. 

83(c). Adding water to the residue from the heated 
marble. 

84. Reaction between calcium carbonate and hydrochloric 
acid. 

85. Reaction between marble and sulphuric acid. 

86. Precipitation of lime soap from permanently hard 
water by a soap solution. Pure castile soap is sodium stear- 
ate, and its formula is NaCi8H 35 2 . Lime soap is calcium 
stearate, and its formula is Ca(C]8H 35 2 ) 2 . Since hard water 
contains calcium sulphate in solution, the other product is 
sodium sulphate. 

87. Preparation of nitric acid from potassium nitrate by 
means of sulphuric acid. 

88. Action of nitric acid on magnesium. Remember that 






FIRST YEAR CHEMISTRY 219 

magnesium is the only -metal that evolves hydrogen from 
nitric acid. 

89(a). Action of nitric acid on copper. The oxide of 
nitrogen evolved was the monoxide, NO. The other prod- 
ucts are copper nitrate and water. 

89(b). Spontaneous oxidation of nitrogen monoxide, NO, 
to nitrogen dioxide, N0 2 , by exposure to air. 

89(c). Reduction of nitrogen monoxide to nitrogen by 
means of phosphorus. 

90. Neutralization of potassium hydrate with nitric acid. 

Experimental work. — Having now considered briefly the 
various theories that have been proposed to explain chem- 
ical action, let us take up again the experimental work of the 
course. We shall take up: first, some experiments to verify 
the laws of chemistry, next some experiments to illustrate 
some of the other points brought out in the theory of chem- 
istry, then a few experiments to determine atomic and molec- 
ular weights, and finally, experiments on bromine, iodine, 
fluorine, arsenic, antimony, and several of the heavy metals. 
These last experiments will consider the elements and their 
compounds in the light of the chemical theory just studied. 
When we say in Experiment 91, and those that immediately 
follow it, that we "verify" a law, it is understood that our 
intention is to show that the law holds for the particular 
case that we try. To verify it completely, i.e., to show that 
the law holds for all cases, would need an endless number 
of experiments. 

Experiment 91. An experiment to verify the Law of Con- 
servation of Mass. Have ready two large beakers, a large 
funnel with filter paper, graduate, horn-pan balance, set of 
smaller weights, brass forceps, tripod, gauze, procelain evap- 
orating dish, ring stand with ring, Bunsen burner, two glass 



220 FIRST YEAR CHEMISTRY 

stirring rods, and a wash-bottle; also some barium nitrate 
and some potassium sulphate. 

Under the Pneumatic Period it was stated that Lavoisier 
formulated the Law of Conservation of Mass, or, as it is 
sometimes called, the Law of Indestructibility of Matter: 
The sum of the weights of the products of a chemical change is 
exactly equal to the sum of the weights of the factors. Let us 
now verify the First Great Law of Chemistry. 

Weigh out exactly 10 grams of thoroly dry barium nitrate 
and dissolve it in about 50 c.c. of distilled water in one of 
the beakers. Then weigh out exactly 7 grams of thoroly 
dry potassium sulphate, and dissolve this in about 50 c.c. 
of distilled water in the other beaker. Bring both solu- 
tions just to a boil, having a stirring rod in each beaker and 
stirring each to hasten the dissolving. Slowly add one so- 
lution to the other, washing the last drops of the solution 
from the first beaker into the second beaker by means of 
che wash-bottle. Note the metathesis. The barium changes 
places with the potassium, and there results barium sul- 
phate, which separates as a white precipitate, and potassi- 
um nitrate, which remains in solution. Write the equation 
for the metathesis which takes place in this experiment. 

Let the sulphate of barium settle while you fit a filter 
paper to the large funnel. Then decant most of the clear 
supernatant liquid into the funnel, catching the filtrate in 
a large beaker. When but little liquid remains with the 
precipitate, shake or stir the contents of the beaker and 
transfer the precipitate to the filter paper, washing out the 
last traces of it from the beaker by means of a stream of 
water from the wash-bottle. Do not lose any of the pre- 
cipitate, and let all the wash-water drain into the original 
filtrate. Boil down the liquid in the beaker to small vol- 
ume and then transfer it to a weighed porcelain evaporating 
dish, rinsing out the beaker with a little distilled water, 



FIRST YEAR CHEMISTRY 



221 



and adding this to the liquid already in the dish. Care- 
fully evaporate it to dryness. If the liquid begins to spat- 
ter toward the end of the heat- 
ing, set the evaporating dish 
into the top of a large beaker 
half full of boiling water, as 
shown in Fig. 72. The heat of 
the steam from the boiling water 
will drive off the last trace of 
moisture from the contents of the 
dish. The apparatus just de- 
scribed is called a steam bath, 
and is often used when there 
must be no loss of substance by 
spattering. When the dish has 
cooled, get the weight of the dish 
and contents. Heat to constant 
weight, and from this calculate 
the weight of the potassium ni- 
trate. Dry the precipitate on the 
filter paper as directed in Experi- 
ment 19. Be sure that you dry 
to constant weight. Dry an un- 
used filter paper to constant 
weight, in order to get the true 
weight of the barium sulphate. 

Finally, compare the sum of the weights of the two 
products with the sum of the weights of the two factors. 

Experiment 92. An experiment to verify the Law of Defi- 
nite Proportions by Weight. Have ready three tall beakers, 
two porcelain evaporating dishes, horn-pan balance, set of 
smaller weights, graduate, tripod, gauze, glass stirring rod, 
two labels, and Bunsen burner; also some sodium carbonate 
and some hydrochloric acid. 




Fig. 72. Steam bath. 



222 FIRST YEAR CHEMISTRY 

Under the Atomic Period it was stated that Proust for- 
mulated the Law of Definite Proportions by Weight: Every 
distinct chemical compound has a fixed and unalterable compo- 
sition. Let us now verify this Second Great Law of Chem- 
istry. 

Weigh out about 20 grams of sodium carbonate; either 
the dry soda or the crystallized salt will answer, but the an- 
hydrous salt is preferable. Dissolve the soda in from 50 
to 60 c.c. of distilled water, heating if necessary. When the 
solution is clear, put exactly 25 c.c. of it in a beaker and 
label it No. 1. Put exactly 25 c.c. of the solution in another 
beaker and label this No. 2. 

To the solution in beaker No. 1, add concentrated hydro- 
chloric acid, slowly, as long as carbon dioxide is evolved. 
Boil down the solution to small volume, transfer it to a 
weighed porcelain evaporating dish and evaporate to dry- 
ness, using the steam bath towards the end of the heating 
when the liquid shows the first signs of spattering. Heat 
to constant weight and get the weight of the residual sodi- 
um chloride. Write the equation for the change that takes 
place when hydrochloric acid is added to sodium carbonate. 

To the solution in beaker No. 2, add hydrochloric acid as 
before, but when effervescence has stopped, add a consid- 
erable excess of the free acid, say about 25 or 30 c.c. Then 
evaporate to dryness, but do this part in the hood. Evapo- 
rate as before to constant weight and get the weight of 
the residual sodium chloride in this case. Compare the 
weight of sodium chloride in this second case with the 
weight of salt obtained when just the proper amount of 
acid was used. 

Experiment 93. Verification of the Law of Multiple Pro- 
portions. No apparatus is necessary for this experiment, 
because the verification of this law consists of calculations 



FIRST YEAR CHEMISTRY 223 

involving the percentage composition of compounds as de- 
termined by more advanced workers than ourselves in chem- 
istry. 

Under the Atomic Period it was stated that from his 
study of the two oxides of carbon and the five oxides of 
nitrogen. Dalton formulated the Law of Multiple Propor- 
tions by Weight: When varying amounts of one sub- 7 
join a fixed quantity of some other, the varying amounts of 
the first bear to each other a simple numerical ratio, such as 
1:2. 1:3. 2:3. etc. Let us now consider Dalton's work in 
detail, and then calculate for ourselves the simple ratios in 
a few cases covered by this Third Great Law of Chemistry. 

A careful quantitative analysis of the oxides of carbon 
showed Dalton that carbon monoxide consists of -12. S6 per 
cent carbon and 57.14 per cent oxygen, while carbon di- 
oxide consists of 27.27 per cent carbon and 72.73 per cent 
oxygen. It is customary to express the composition of 
compounds in percentage form, but the figures given above 
do not reveal any simple ratio between the amounts of oxy- 
gen. Suppose, however, that the same amount of carbon, 
say one gram, was started with in two different experi- 
ments, and that one gram of carbon was oxidized to the 
monoxide, while the other gram of carbon was oxidized to 
the dioxide. The amount of oxygen necessary in the first 
case may be found from the percentage composition of the 
monoxide by means of the following proportion: The per 
cent of carbon in carbon monoxide is to the per cent of 
oxygen in carbon monoxide as the gram weight of the carbon 
is to the gram weight of the oxygen. L'sing the figures, 
this proportion becomes 42.86:57.14= l:x. The calcula- 
tion of x amounts to dividing 57.14 by 42.^6: this division 
shows that one gram of carbon uses 1.33 grams of oxygen 
when forming the monoxide. In like manner, by dividing 
72.73 by 27.27 we find that 2.66 grams of oxygen unite with 



2^4 



FIRST YEAR CHEMISTRY 



one gram of carbon to form the dioxide. Now that we 
have adopted a constant weight for one element we see 
easily that the ratio between the varying amounts of the 
other element is 1.33:2.66, or 1:2. These figures are re- 
peated for emphasis in the following table: 



Oxides of Carbon 


Name of 
substance 


Percentage 
composition 


Gram Weights 


Simple 

ratio 

between 

the weights 

of oxygen 


Carbon 


Oxygen 


Unit 

weight of 

carbon 


Weight of 

oxygen 

that unites 

with unit 

weight of 

carbon 


Carbon monoxide 
Carbon dioxide 


42.86% 
27.27% 


57.14% 
72.73% 


1 g- 
1 g- 


1.33 g. 
2.66 g. 


1:2 



Let us now consider the five oxides of nitrogen. The 
table given below shows the percentage composition of 
each of the five oxides of nitrogen. The next to the last 
column shows the amounts of oxygen that combine with 
unit weight of nitrogen to form the different oxides, these 
amounts having been obtained from the percentages of the 
elements in the corresponding oxides. The last column 
shows the ratio between the oxygen in nitrous oxide and 
the oxygen in each of the other oxides. 

The table shows that the ratio between the oxygens in the 
first two oxides of nitrogen is 1:2, that between the oxy- 
gens in the first and third oxides is 1:3, and so on down 
the list. All this may be condensed into the statement that 
the ratio between the varying amounts of oxygen in the 
five oxides of nitrogen is 1:2:3:4:5. 

Let us now calculate the simple ratio between the oxy- 



FIRST YEAR CHEMISTRY 



226 



gens in the two oxides of sodium, between the chlorines in 
the two chlorides of iron, and between the hydrogens in 
the two substances defiant gas and marsh gas, — two gase- 



Oxides of Nitrogen 


Name of 
substance 


Percentage 
composition 


Gram Weights 


Simple 

ratio 

between 

the weights 

of o xy gen 


Nitrogen 


Oxygen 


Unit 
weight of 
nitrogen 


Weight of 

oxygen 

that unites 

with unit 

weight of 

nitrogen 


Nitrous 

oxide 
Nitric 

oxide 
Nitrogen 

tnoxide 
Nitrogen 

peroxide 
Nitrogen 

pentoxide 


63.6 
46 6 
36.8 
30.4 
25.9 


36.4 
53.4 
63.2 
69.6 
74.1 


1 
1 
1 
1 

1 


0.56 
1.11 
1.71 
2.29 

2.86 


1:2 
1:3 
1:4 
1:5 



ous hydrogen compounds of carbon that Dalton studied. 
The necessary data are given in Problems 93(a), 93(b), and 
93(c). Arrange the data there given together with the re- 
sults of the calculations in the form of tables as in the case 
of the oxides of nitrogen. Make a table for each problem 
and let both the tables and the mathematical work appear 
on the proper left-hand page in the laboratory note- 
book. 

Problem 93(a). — The two oxides of sodium have the fol- 
owing composition: Sodium monoxide contains 74.19% of 
sodium and 25.81% of oxygen; sodium dioxide contains 
58.97% of sodium and 41.03% of oxygen. Calculate the 
simple ratio between the varying amounts of oxygen that 
unite with unit weight of sodium. Ans. 1 : 2. 



226 FIRST YEAR CHEMISTRY 

Problem 93(b). — The two chlorides of iron have the fol- 
lowing composition: Ferrous chloride contains 44.1% of 
iron and 55.9% of chlorine; ferric chloride contains 34.4% 
of iron and 65.6% of chlorine. Calculate the simple ratio 
between the varying amounts of chlorine that unite with 
unit weight of iron. Ans. 2 : 3. 

Problem 93(c). — Of the many gases containing carbon 
and hydrogen, marsh gas and olefiant gas are the more 
common. These two gases have the following composition: 
Olefiant gas contains 85.7% of carbon and 14.3% of hydro- 
gen. Marsh gas contains 75.0% of carbon and 25.0% of hy- 
drogen. Calculate the simple ratio between the varying 
amouns of hydrogen that unite with unit weight of carbon 
to form these two gases. Ans. 1 : 2. 

Experimen 94. Verification of the Law of Definite Pro- 
portions by Volume. No apparatus is necessary for this ex- 
periment because the verification of this law consists of the 
examination of the results of experiments already performed. 
These results will be taken partly from our own work this year 
and partly from the work of more advanced experimenters. 

Under the Atomic Period it was stated that Gay-Lussac 
formulated the Law of Definite Proportions by Volume: 
In any chemical change the relative volumes of the gaseous 
factors and products bear to each other a simple numerical 
ratio, such as 1:2, 1:3, 2:3, etc.; and incidentally, this 
ratio is always expressed by the number of molecules of 
the gaseous factors and products in the written equation. 
Let us now verify this Fourth Great Law of Chemistry in 
a number of cases, and then see how it can serve us in de- 
termining, without actual experiment, the volumes of the 
gaseous substances that figure in any chemical change. 

We have done several experiments in which gaseous terms 
figured. The first experiment that contained more than 



FIRST YEAR CHEMISTRY 227 

one gaseous term was the experiment on the decomposition 
of water by means of electricity. In that case we obtained 
two volumes of hydrogen and one volume of oxygen. In 
the equation for the decomposition, 

2H 2 = 2H 2 +0 2 , 

we notice that the coefficients of the two gaseous terms, 
hydrogen and oxygen, are 2 and 1 ; we notice also that these 
coefficients are identical with the relative volumes of these 
two gases as caught in the tubes of the electrolysis appara- 
tus. The simple ratio in this case is 2:1. This coinci- 
dence does not apply to the coefficient of the H 2 0, for that 
substance is not a gas; this happy coincidence relates only 
to gaseous terms, and it has been shown by experiment that 
it holds good for all changes involving gaseous terms. 

Another experiment was the burning of sulphur in oxy- 
gen. If we had measured the volumes of the gases care- 
fully we would have found that one volume of oxygen pro- 
duced just one volume of sulphur dioxide. Note that in the 
equation, 

S + 2 = S0 2 , 

the coefficients of the gaseous terms oxygen and sulphur di- 
oxide are identical with the volumes as given above. The 
ratio in this case is 1 : 1. 

In making the second oxide of sulphur two volumes of 
sulphur dioxide and one volume of oxygen produce two vol- 
umes of sulphur trioxide when the product is in the gas- 
eous state. Note again that in the equation, 

2S0 2 +0 2 =2S0 3 , 

the coefficients of the terms are identical with the volumes 
as given above. Since all three terms of this equation are 
gases the ratio is 2 : 1 : 2. 



228 FIRST YEAR CHEMISTRY 

Other experiments we have performed include the follow- 
ing: 

When hydrogen is burned in oxygen, if the volume of 
the resulting steam is measured, we find that two volumes 
of hydrogen and one volume of oxygen produce two vol- 
umes of water vapor as represented by the equation, 

2H 2 +0 2 =2H 2 0. 

In this case the simple ratio is 2 : 1 : 2. 

One volume of hydrogen produces one volume of hydro- 
gen sulphide as represented by the equation, 

H 2 + S = H 2 S, 

In this case the simple ratio between the volumes of hydro- 
gen and hydrogen sulphide is 1 : 1. 

One volume of oxygen produces one volume of carbon di- 
oxide as represented by the equation, 

C + 2 = C0 2 . 

In this case the simple ratio between the oxygen and the 
carbon dioxide is 1 : 1. 

This principle in regard to gas volumes may be used in 
another way. If the equation for a certain reaction is 
known, the volumes of the gaseous factors and products 
may be determined directly from the equation instead of 
by actual measurement. For instance, the coefficients in 
the last equation — for the oxidation of carbon — show directly 
that one volume of oxygen produces one volume of carbon 
dioxide. 

Write in the notebook the equations indicated below, and 
tell from the equations the volumes of the gaseous terms 
figuring therein. 

(1) Reduction of steam to hydrogen by means of mag- 
nesium. 



FIRST YEAR CHEMISTRY 229 

(2) Complete oxidation of hydrogen sulphide. 

(3) Reducing carbon dioxide, C0 2 , to carbon monoxide, 
CO, by means of hot zinc. 

(4) Oxidizing carbon monoxide to carbon dioxide. 

(5) Reduction of carbon dioxide to carbon monoxide by 
means of hot carbon. 

(6) Burning hydrogen in chlorine. 

(7) Spontaneous oxidation of nitrogen monoxide, NO, 
to nitrogen dioxide, N0 2 , by exposure to air. 

(8) Reduction of nitrogen monoxide to nitrogen by 
means of phosphorus. 

Experiment 95. An experiment to verify the Law of Boyle. 

(Lecture Experiment.) Have ready a hard glass tube of 7 
or 8 mm. bore and at least a meter long, triangular file, 
small beaker, large porcelain mortar, ring stand and clamp, 
meter rod, and a Boyle tube with stand to support it; also 
about 500 grams of clean, dry mercury. 

It is well known that air occupies space and exerts pres- 
sure. Numerous examples might be cited, e.g., the action 
of a bicycle pump, and the resistance noticed when running 
against a strong wind. It may not be so well known that 
when a gas is subjected to pressure its volume diminishes, 
and that this change may be expressed mathematically. 
Under the Period of Robert Boyle it was stated that Boyle 
formulated the law for the effect of pressure on the volume 
of a gas, this law being now called the Law of Boyle: The 
volume of a gas is inversely proportional to the pressure to 
which it is subjected. Let us now verify this law by meas- 
uring the volume of a definite amount of confined air under 
different pressures. 

Making a crude barometer. — Let us first make a crude 
barometer, since this instrument must be used later in this 
experiment. Select a clean and dry hard glass tube of 7 



230 



FIRST YEAR CHEMISTRY 



or 8 mm. bore and at least a meter long. 
Heat the end in the blast lamp, and 
close it without forming a large bulb at 
the end. Fill this tube to within 2 or 3 
cm. of the open end with clean, dry mer- 
cury. Put the finger over the open end, 
invert the tube and allow the inclosed 
air to travel slowly to the other end of 
the tube and back, gathering in the small 
bubbles of air that cling to the sides of 
the tube. Now fill the tube completely 
with mercury, place the finger over the 
end, and invert the tube over a mortar 
half full of mercury. When the open 
end of the tube is well immersed, re- 
move the finger, bring the tube into 
upright position, and clamp it thus 
with a clamp to the ring stand, taking 
the precaution that the end of the tube 
is under the surface of the mercury in 
the mortar, but that it does not touch 
the bottom of the mortar. Note that 
the mercury column falls and soon comes 
to rest, leaving an empty space in the 
top of the tube. This space is called Torri- 
celli's Vacuum, after the man who first 
did this experiment. See Figure 73. 
With a meter stick measure the hight of 
the column of mercury above the sur- 
face of the mercury in the mortar, by 
putting a meter stick alongside of the 
glass tube with the end of the stick rest- 
ing on the surface of the mercury in the 
mortar. Read the hight to millimeters. This reading should 



Fig. 73. A crude 
barometer. 



FIRST YEAR CHEMISTRY 231 

be somewhere between 700 and 800 mm.; it varies with the 
conditions of the atmosphere. Record the reading in the 
notebook. 

The pressure of the air. — Let us now see how this hight 
is used to measure the pressure of the air. Since there is 
a vacuum in the top of the tube there is no pressure on the 
upper surface of the mercury column. This column of mer- 
cury is kept up in the tube by means of the weight of the 
air on the surface of the mercury in the mortar. It has 
been found by experiment that the size of the tube does 
not affect the hight of the mercury column. Suppose, for 
convenience, that the cross section of the tube is exactly 
one square centimeter. The volume of the mercury col- 
umn contains, then, as many cubic centimeters as the col- 
umn is linear centimeters high. If this column were com- 
posed of water instead of mercury, the number expressing 
its volume would also express its weight in grams, since 
one cubic centimeter of water weighs one gram. Mercury 
is 13.6 times as heavy as water. Therefore, the weight of 
the mercury column may be found by multiplying the 
number expressing its volume by 13.6, and this weight 
represents the pressure of the air on one square centimeter. 
Using the equivalents 2.54 cm. =1 in. and 28 grams = 1 
ounce, calculate (a) the pressure in grams per square inch, 
(b) the pressure in ounces per square inch, and (c) the pres- 
sure in pounds and ounces per square inch. The final re- 
sult should come out about 15 pounds. In ordinary prac- 
tice the hight of the mercury column is used rather than 
the more correct, tho not so easily determined, weight of mer- 
cury column. All these calculations should be made on 
the left-hand page in the notebook. 

Work with the Boyle tube. — Let us now proceed to the 
experiment proper. Have ready a Boyle tube — a glass tube 
of 8 to 10 min. bore, of uniform caliber, about 1.5 meters 



232 



FIRST YEAR CHEMISTRY 



long, closed at one end, and bent to form two parallel arms, 
one of which (the open one) is at least three times as long 

as the other. If this tube is not 
already attached to a support, 
fasten it securely in an upright 
position to a large ring stand or 
other suitable support. 

Pour enough mercury into the 
tube to seal the bend, i.e., to fill 
the bottom of the U, and to ex- 
tend a short distance up into the 
parallel arms. Tip the tube 
around in different directions so 
as to make the mercury stand 
at the same hight (A and A' in 
Figure 74) in both arms when 
the tube is upright. The column 
of air, AC, in the confined shorter 
arm is under the same pressure 
as the air outside the tube; it is 
then said to be under a pressure 
of one atmosphere. From this 
time on avoid as far as possible 
any heating of the confined air 
from contact with the hands or 
breath. Measure the length AC 
of the column of confined air. 

Now let us increase the pres- 
sure on the inclosed air to two 

atmospheres, by adding mer- 
Fig. 74. The Boyle tube. cury tQ the Qpen end m the 

distance from the level B to the level B' is equal to the 
hight of the barometer as determined in the early part of 
this experiment. Why does it not double the pressure on 




FIRST YEAR CHEMISTRY 233 

the confined air simply to add a column of mercury to the 
long arm equal in length to the hight of the barometer? 
Now measure the length BC of the column of confined air, 
and compare it with its original length. What effect has 
doubling the pressure on the air had on its volume? The 
law holds good for all pressures; tripling the pressure re- 
duces the volume to one third its original volume, and halv- 
ing the pressure doubles the volume. If P and V stand for 
the original pressure and volume, and P' and V for the 
new pressure and volume, the law may be expressed by the 
following proportion: V:V'=P':P; this proportion is used in 
calculating the volume that a gas would assume under a 
new pressure. Occasionally the proportion is expressed in 
equation form thus: PV = P'V; but this form is not so 
convenient, for our use, as the proportion. 

Application of the Law of Boyle. — The proportion given in 
the preceding paragraph may be used in such a problem as 
the following : A gas measures 200 c.c. when under a -pres- 
sure of 740 mm. What will its volume become if the pressure 
changes to 760 mm.? This problem should be solved as fol- 
lows: First write the formula V:V'=P':P. Substitute for 
V the volume given in the problem, i.e., 200 c.c, also sub- 
stitute for P the pressure at which the gas measures 200 c.c, 
i.e., 740 mm. Let V be the unknown volume; then P' should 
be replaced by 760 mm. The rest of the solution consists, of 
course, of multiplying the extremes and then multiply the 
means, and finally dividing the product of the extremes by 
the coefficient of V. The actual work should look like this: 

V : V = P' : P 

200 : V = 760 : 740 

760 V'= 200X740 

200X740 



760 
V'= 194.7 c.c 



234 FIRST YEAR CHEMISTRY 

When the answer has been obtained its probable accuracy 
should be checked as follows: The pressure changes from 
740 mm. to 760 mm. Does the pressure increase or dimin- 
ish? It increases. What effect does this have on the vol- 
ume? Increasing the pressure diminishes the volume. The 
answer agrees with this in that it is smaller than the vol- 
ume started with. This shows that the general formula was 
properly used. 

Problems involving the Law of Boyle. — Do the following 
problems on reducing gas volumes, using the general for- 
mula just illustrated. Record all the calculation on the 
left-hand page of the notebook and carry each answer out 
to one decimal place. 

95(a). A gas measures 900 c.c. at 745 mm. pressure. 
Find its volume at 760 mm. pressure. Ans. 882.2 c.c. 

95(b). A gas measures 400 c.c. at 775 mm. pressure. 
Find its volume at 760 mm. pressure. Ans. 407.8 c.c. 

95(c). A gas measures 500 c.c. at 750 mm. pressure. 
Find its volume at 770 mm. pressure. Ans. 487.0 c.c. 

95(d). A gas measures 700 c.c. at 770 mm. pressure. 
Find its volume at 750 mm. pressure. Ans. 718.6 c.c. 

95(e). A gas measures 200 c.c. at 760 mm. pressure. 
Find its volume at 770 mm. pressure. Ans. 197.4 c.c. 

95(f). A gas measures 300 c.c. at 760 mm. pressure. 
Find its volume at 750 mm. pressure. Ans. 304.0 c.c. 

The Barometer. — A high-grade barometer is similar in principle to 
the crude barometer that was made in the experiment on the Law 
of Boyle. The crude barometer was simply a glass tube, closed at one 
end, filled with mercury, and inserted in a mortar of mercury. See 
Fig. 73. The hight of the column of mercury was measured by put- 
ting a meter stick alongside of the glass tube with the end of the stick 
resting on the surface of the mercury in the mortar. This method 
naturally does not give an accurate reading. The main errors are not 
obviated by fastening the meter stick rigidly to the tube, because the 
level of the mercury in the mortar changes slightly as the mercury 



FIRST YEAR CHEMISTRY 



235 



column rises or falls with the increase or de- 
crease in the pressure of the atmosphere. 

In a high-grade barometer the zero point of 
the measuring scale is indicated by the point 
of an ivory peg. A, as shown in Fig. 75. The 
bottom of the cistern of mercury is movable, 
and may be raised or lowered by means of the 
screw B, till the surface of the mercury just 
touches the point of the ivory peg. The cis- 
tern is often covered to keep out the dust, 
but this cover is not hermetically sealed; hence, 
the surface of the mercury in the cistern is 
subjected to the full pressure of the atmos- 
phere. The long glass tube is often inclosed 
in a metallic tube for protection. The metallic 
tube sometimes carries a small thermometer, 
but in our work this may be neglected. The 
top of the mercury column may be seen thru 
the hole C in the metallic collar D; this collar 
is graduated in millimeters on the left-hand 
side and in inches on the right-hand side. 
The sliding scale E may be raised or lowered 
by means of the screw F. 

The sliding scale E carries a vernier, i.e., 
a small scale made by taking a length equal 
to 9 mm. and dividing it into 10 equal parts. 
The object of the vernier is to facilitate 
reading the instrument to tenths of a milli- 
meter. Its use is illustrated by Fig. 76. Sup- 
pose the mercury stands a little above 749 mm. 
but under 750 mm. It is difficult to estimate 
to tenths of a millimeter by the eye alone. 
The graduations on the vernier are of such a 
nature that when the surface of the mercury 
stands between any two marks on the milli- 
meter scale it is always possible after bringing 
the vernier scale down so that its lower edge 
coincides with the top of the mercury column, 
to find some mark on the vernier that coin- 
eidea with a mark on the millimeter scale. 
The number of the line on the vernier that co- 
incides thus with some line on the millimeter 
scale indicates the number of tenths of a milli- 



Fig. 



75. A high grade 
barometer. 



236 



FIRST YEAR CHEMISTRY 



meter that must be added to the last full millimeter reading below 

the surface of the mercury. The correct reading in Fig. 76 is, 

therefore, 749.5 millimeters. 

When reading the barometer, one must observe the following details: 
(1) Set the mercury in the cistern to the zero point by turning 

the screw at the bottom till the surface of mercury just touches the 

point of the ivory peg. 

(2) Set the sliding scale so 
that its lower edge coincides 
with the top of the mercury 
column. 

(3) Read the hight, — mil- 
limeters from the scale on 
the metallic collar, and tenths 
of a millimeter from the ver- 
nier scale. 

Experiment 96. An ex- 
periment to verify the 
Law of Dalton. (Lecture 
Experiment.) Have 
ready a copper boiler, 
Bunsen burner, ring 
stand and clamp, large 
battery jar, large beaker, 
250 c.c. flask with one- 
holed rubber stopper, 
tubing clamp, platform 
balance, set of iron 
weights, glass tubing, 
rubber tubing, and ice . 

Under the Atomic Pe- 
riod it was stated that 
Dalton formulated the 
Law of Expansion of 
Gases by Heat: The 
volume of a gas increases or decreases by one two hundred and 




Fig. 76. The vernier. 



FIRST YEAR CHEMISTRY 237 

seventy-third of its volume as measured at 0° Centigrade for 
every degree rise or fall in temperature. Sometimes the law 
is stated as follows: The volume of a gas varies directly as 
the temperature on the "absolute sealed Let us now verify 
this law by measuring the volume of a certain amount of 
air at 100°C, then its volume after it has been cooled to 0°C. 

First crush enough ice to fill a large battery jar. 

To a dry 250 c.c. flask fit a one-hole rubber stopper, hav- 
ing a piece of glass tubing flush with the under side of the 
stopper and protruding about 2 cm. above. To this tube 
fit a piece of rubber tubing about 25 cm. long, carrying a 
tubing clamp at the end farthest from the stopper. 

Make sure that the flask and fittings are perfectly dry; get 
the weight of flask and fittings, using the platform balance; 
clamp the flask and fittings by the stopper (not by the neck 
of the flask), and place it as deep as possible in the copper 
boiler. Fill the boiler with water to within an inch of the 
top, and set the flame underneath, being sure that the tub- 
ing clamp is open. 

When the water has begun to boil, keep it boiling for three 
or four minutes to make sure that the air in the flask is at 
100°C. Then close the tubing clamp and remove the flask 
from the boiler. Let it cool until it is about hand warm; 
then pack it in crushed ice in a battery jar, covering it com- 
pletely with ice. Place a large beaker of ice-water upon the 
ice in the jar and put the end of the rubber tubing into the 
ice-water. Open the tubing clamp and allow water to run 
into the flask, being careful always to keep the end of the 
tubing below the surface of the water in the beaker. 

Let the flask remain in the ice five minutes, so that the 
temperature of the air in the flask may become 0°C. 

Some water will be caught in the tubing; this must be 
saved. Compress the end of the rubber tubing with the 
fingers, raise it to a vertical position, loosen the stopper and 



238 FIRST YEAR CHEMISTRY 

allow the water in the tubing to fall into the flask. Re- 
move the flask and fittings from the ice, wipe it thoroly dry 
on the outside and get the weight of flask, fittings and water. 
How many c.c. of water ran in? By how many c.c., then> 
did the air in the flask contract in cooling from 100°C. 
to 0°C? 

Fill the flask to the brim with water and insert the stop- 
per, forcing the water up the tubing. Make sure the rub- 
ber tubing is full of water, adding a few drops from the 
wash-bottle if necessary; close the tubing clamp as before. 
Wipe dry on the outside and weigh again. How many 
grams of water in the flask and tubes? What, then, is the 
capacity of the flask and tubes? 

From the data thus prepared calculate: 

(a) The volume of air experimented on at 100°C, i.e., 
the volume of the flask. 

(b) The contraction in this volume in cooling from 100°C. 
to 0°C. 

(c) The volume at 0°C. 

(d) The amount the volume at 0°C. would expand from 
0°C to 100°C. 

(e) The amount it would expand from 0°C. to 1°C. 

(f) The amount 1 c.c. of air would expand from 0°C. to 
1°C. 

This last numerical value is called the coefficient of ex- 
pansion, and should come out about 0.00366 or about ¥ ^ F . 
It holds good for all gases. 

If one c.c. of gas measured at 0°C. gains or loses •$$■$ of 
its volume for every degree heated or cooled respectively, 
at what temperature does it become 2 c.c. ? What would be 
the volume of 1 c.c. of gas measured at 0°C. if cooled to 
— 1°C. ? What would it seem to become if cooled to — 273°C. ? 
This point, — 273°C, is called the absolute zero and has never 
been obtained ; however, — 250°C. has been reached by means 



FIRST YEAR CHEMISTRY 



239 



273° 



73° 



<f~ 



C 

-100° 



0° 



of liquified gases. A scale having degrees the same size as 
centigrade degrees with its zero mark 273° below 0°C, its 
273 mark at 0°C. and its 373° mark at 100°C. is called the 
absolute scale. See Fig. 77. Tempera- 
ture measured on the absolute scale is 
called absolute temperature. The abso-373 -- 
lute temperature of a gas is always de- 
termined by adding 273 to the centi- 
grade temperature of the gas. 

If V and t stand for the original vol- 
ume and temperature, and V and t' for 
the new volume and temperature, the 
law may be expressed by the following 
proportion: V : V = 273 + t : 273 -ft'; 
this proportion is used in calculating 173°- • 
the volume a gas would assume under a 
new temperature. 

Application of the Law of Dalton. — The 
proportion given in the preceding para- 
graph may be used in such a problem as 
the following: A gas measures 200 c.c. 
when at a temperature of 27°C. What 
will its volume become if the temperature 
changes to 0°C.f This problem should 
be solved as follows: First write the 
formula V : V'=273+t : 273 +t'. Sub- 
stitute for V the volume given in the problem, i.e., 200 c.c; 
also substitute for t the temperature at which the gas meas- 
ures 200 c.c, i.e., 27°C. Let V be the unknown volume; 
then V should be replaced by 0°. The rest of the solution 
consists, of course, of multiplying the extremes, then multi- 
plying the means, and finally dividing the product of the ex- 
tremes by the coefficient of V'. The actual work should 
look like thi^: 



100° 



-200° 



•273 ( 



Fig. 77. The abso- 
lute scale. 



240 FIRST YEAR CHEMISTRY 

V : V'= 273 +t : 273 +t' 
200 : V'=273+27 : 273+0 
200 : V = 300 : 273 

300 V'= 200X273 
200X273 



300 
V' = 182.0 c.c. 

When the answer has been obtained its probable accuracy 
should be checked as follows: The temperature changes 
from 27° to 0°. Does the temperature increase or decrease ? 
It decreases. What effect does this have on the volume? 
Decreasing the temperature decreases the volume. The an- 
swer agrees with this in that it is smaller than the volume 
started with. This shows that the general formula was 
properly used. 

Problems involving the Law of Dalton. — In the following 
problems on reducing gas volumes use the general formula 
just illustrated. Record all the calculations on the left- 
hand page of the notebook, and carry each answer out to 
one decimal place. Also check the probable accuracy of 
the answer. 

96(a) A gas measures 400 c.c. at 27°C. Find its volume 
at 0°C. Ans. 364.0 c.c. 

96(b) A gas measures 900 c.c. at 0°C. Find its volume 
at 27°C. Ans. 989.0 c.c. 

96(c) A gas measures 500 c.c. at 20°C. Find its volume 
at 10°C. Ans. 482.9 c.c. 

96(d) A gas measures 100 c.c. at — 73°C. Find its vol- 
ume at 0°C. Ans. 136.5 c.c. 

Application of both the Law of Boyle and the Law of Dal- 
ton. — It is often necesary to reduce the volume of a gas for 
changes in both pressure and temperature. Such a problem 
might read as follows: A gas measures 200 c.c. at 780 mm, 



FIRST YEAR CHEMISTRY 241 

pressure and a temperature cf 27°C. What will its volume 
become at 760 mm. pressure and 0°CJ In such a problem find 
out first what the volume would become if only the pres- 
sure changed. Consider this answer the new known volume 
and find out what this w T ould become when subjected to the 
change in temperature indicated in the problem. The work 
would be as follows: 



V : V 


= P' : P 


V 


: V 


= 273+t : 273 +t' 


200 : V 


= 760 : 780 


205.2 


: V 


= 273+27 : 273+0 


760 V 


= 200X780 


205.2 


: V 


= 300 : 273 


V 


200X780 
760 




V 


205.2X273 
300 


V 


= 205.2 c.c. 




V' 


= 186.7 c.c. 



Problem involving both the Law of Boyle and the Law of 
Dalton. — In the following problem reduce the volume of the 
gas first, according to the Law of Boyle, and then, consid- 
ering the answer obtained from this as the known volume, 
reduce this according to the Law of Dalton. Check the 
probable accuracy of the answer. 

96(e) A gas measures 500 c.c. at 770 mm. pressure and 
27°C. Find its volume at 760 mm. pressure and 0°C. 

Ans. 460.9 c.c. 

Experiment 97. Weight and specific gravity of air. (Lect- 
ure Experiment.) Have ready a large prescription bottle 
of two to three liters capacity and fitted with a one-holed 
rubber stopper, platform balance, set of large weights, set 
of smaller weights, a good aspirator, tubing clamp, a piece 
of glass tubing about 10 or 15 cm. long, a piece of heavy 
rubber tubing about 10 or 15 cm. long, and the pneumatic 
trough; also some vaseline. 

Under the Pneumatic Period it was pointed out that 
Cavendish saw the necessity of obtaining the exact specific 



242 FIRST YEAR CHEMISTRY 

gravities of gases in order to distinguish one gas from an- 
other. Since the specific gravity of a gas is generally re- 
ferred to air as a standard, it will be of advantage to us to 
determine the exact weight of air and its specific gravity 
referred to water, which is the standard for liquids and solids. 

Make sure that the bottle is clean and perfectly dry. Thru 
the stopper, which should fit tightly, pass a piece of tightly 
fitting glass tubing about 10 or 15 cm. long. Slip over the 
outer end of the glass tube a piece of thick walled rubber 
tubing about 10 or 15 cm. long. Put the tubing clamp on 
the rubber tube. If you do not feel sure that the fittings 
are air-tight, moisten with vaseline those surfaces of the 
rubber stopper and tube that are in contact with the glass. 

Fill the pneumatic trough with water to within two or 
three centimeters of the top. 

Insert the stopper and fittings, close the tubing clamp, 
and weigh the bottle and fittings full of air accurately to 
centigrams. Record the weight in the notebook. Do not 
remove the weights from the balance, because the differ- 
ence between this weight and the next one will be very 
small, and this difference must be determined as accurately 
as possible. Make sure that the end of the aspirator to 
which the tube is to be attached is perfectly dry. Attach 
the bottle to the aspirator and turn on the water to get the 
suction. Do not turn off the water until after the bottle has 
been disconnected from the aspirator. Open the tubing clamp 
and suck air from the bottle for about five minutes. Close 
the tubing clamp, remove the bottle and fittings from the 
aspirator, and set it back on the balance. To get the weight 
of air removed, put small brass weights on the pan beside 
the bottle till equilibrium is established. Record the weight 
in the notebook. 

To determine the volume of air removed from the bottle, 
immerse the bottle inverted in the pneumatic trough. With 



FIRST YEAR CHEMISTRY 243 

all the fittings under water, open the tubing clamp and let 
water run into the bottle. What causes the water to rush 
in? Keep the fittings still underwater and raise or lower 
the bottle till the water inside of it is on the same level 
with the water on the outside of it. In order to get the 
levels the same it will probably be necessary to tip the bottle 
considerably, but this tipping will have no effect on the 
amount of water caught in the bottle. Then close the tub- 
ing clamp and remove the bottle from its bath. Wipe the 
bottle and fittings dry on the outside and weigh it on the 
platform balance. Get within one gram of the true weight. 
Greater accuracy at this point is unnecessary. Record the 
weight in the notebook. 

From the weight of the bottle with the water that ran in, 
and the weight of the bottle empty, calculate the weight of 
water that ran in. Remembering that one gram of water 
occupies 1 c.c, get the volume of the water that ran in, i.e., 
the volume of air that was pumped out. You have already 
determined the weight of the air removed. From the above 
data, get the weight of 1 c.c. of air in the room at the time 
of doing the experiment. 

We have now determined the weight of unit volume of 
air, i.e., we have determined the density of air. The spe- 
cific gravity of a substance is the number of times heavier or 
lighter unit volume of that substance is than unit volume of 
some substance taken as a standard. At the opening of this 
experiment it was stated that water was often used as a 
standard for liquids and solids. One c.c. of water weighs 
one gram. From this, and from the weight of 1 c.c. of air 
as you found it, calculate the specific gravity of air referred 
to water, by dividing the weight of 1 c.c. of air by the weight 
of 1 c.c. of water. 

The density of air is roughly about 0.0013 gram, but it varies 
slightly from day to day according to atmospheric conditions. 



244 FIRST YEAR CHEMISTRY 

Experiment 98. Weight and specific gravity of carbon 
dioxide. Have ready a 500 c.c. flask fitted with a one-hole 
rubber stopper, tubing clamp, glass tubing, rubber tubing, 
gas balance, platform balance, set of large weights, set 
of special weights to use with the gas balance, ther- 
mometer and barometer; also a bag of carbon dioxide. 

In determining the specific gravity of a gas, air more fre- 
quently than water is taken as a standard. In the preced- 
ing experiment we found the weight of 1 c.c. of air. In this 
experiment we shall find the weight of 1 c.c. of carbon di- 
oxide. A comparison of the two weights will give us the 
specific gravity of carbon dioxide referred to air. This 
comparison must, however, be made at what are called 
standard conditions, i.e., at a temperature of 0°C. and under 
760 mm. pressure. The expression " standard conditions" 
is often written by S. T. P., — the abbreviation for Stand- 
ard Temperature and Pressure. 

Make sure that the flask is clean and perfectly dry. Fit 
it with a one-hole rubber stopper, thru which passes a tightly 
fitting glass tube long enough to reach to the bottom of the 
flask and to project a couple of centimeters above the top 
of the stopper. To the outer end of the glass tube attach 
a short piece of rubber tubing carrying a tubing clamp. The 
flask is now ready for weighing on the gas balance. 

Directions for using the gas balance. See Fig. 78. 

Before weighing anything, see that the pointer comes to rest oppo- 
site the central mark on the black graduated scale. If it does not 
come to rest at the central mark, turn the screw at the right of the 
beam at the top of the balance till the pointer does come to rest at the 
central mark. 

To stop the pointer from swinging, press the thumb knob at the 
middle of the black piece that carries the scale till the top edge of 
the scale touches the pointer. 

Put the article to be weighed in the left-hand pan. Put the weights 
in the right-hand pan. 



FIRST YEAR CHEMISTRY 



245 



Use only the special weights in the boxes in front of the balance. 
For making weighings of ten grams or over use the weights from the 
box in front of the balance. For making weighings under ten grams 
use the rider on the beam at the top of the balance. 

The graduations 
on the beam marked 
1. 2. 3, etc., indicate 
grams. Each gram 
division is graduated 
to tenths. Read 
all weighings with 
this balance to 
tenths of a gram at 
least. Read closer 
if possible, i.e., to 
hundredths of a 
gram by estimating 
the tenths of the 
smallest division on 
the beam. 

Set the flask 
and fittings with 
the tubing clamp 
open on the gas 
balance, and get 
its exact weight. 
Do not remove 
the weights from 
the pan, for the 

increase in weight 

, it Fig. 78. The gas balance, 

due to carbon di- 
oxide will be very small and this increase must be determined 
as accurately as possible. 

Remove the flask from the balance and loosen the stopper. 
Attach a rubber bag full of carbon dioxide to the fittings 
of the flask, and pass the gas slowly thru the fittings into 
the flask till a match held where the stopper is loosened is 




246 FIRST YEAR CHEMISTRY 

extinguished. Insert the stopper, close the tubing clamp 
and remove the rubber bag. Open the tubing clamp for an 
instant to let out any excess of gas and then close it again. 
Replace the flask on the gas balance and find the gain 
in weight of the flask and contents. To make sure the flask 
is full of carbon dioxide loosen the stopper, pass in some 
more gas and weigh again. If there is a further gain in weight 
continue alternately passing in gas and weighing till you 
have attained constant weight. When the weight is finally 
constant, read the barometer and the thermometer. 

Next fill the flask with water and insert the fittings so 
that the water fills them to the top. Close the tubing clamp 
and wipe the flask and fittings dry on the outside. Get the 
weight of the flask and fittings full of water on the platform 
balance. So much for the experimental work; you are now 
ready for the arithmetical work which, by the way, must 
all be done on the left-hand pages of the notebook. 

From the weight of the flask and fittings empty and the 
weight when full of water, calculate the weight of the water 
that fills the flask. What, then, was the volume of the 
water, i.e., what was the volume of air in the flask? If you 
knew the weight of 1 c.c. of air at the conditions under 
which you were working, it would be an easy matter to 
multiply this weight by the capacity of the flask and to get 
the weight of air in the flask. Since the weight of air changes 
with the temperature and the pressure, it is more convenient 
to reduce the volume of air as you measured it to its volume 
at S. T. P., and then multiply it by 0.00129 gram, which 
is the weight of 1 c.c. of air at S. T. P. 

To reduce to S. T. P., use the method described in the 
problem involving both the Law of Boyle and the Law of 
Dalton at the end of Experiment 96, i.e., find out first, 
what the volume of the air in the flask would become at 760 
mm. pressure, using the proportion V : V' = P' : P; for the 



FIRST YEAR CHEMISTRY 247 

present you neglect the temperature. When you have de- 
termined the volume at 760 mm. pressure, find out what 
that new volume would become if changed from the observed 
temperature to 0°C. Here, of course, you use the propor- 
tion V : Y'=273+t ■ 273 +t'. When you have found the 
volume at S. T. P., multiply it by 0.00129 and you have the 
weight of air in the flask at the observed temperature and 
pressure. 

The gain in weight noted when the carbon dioxide was put 
into the flask plus the weight of the air the flask held is the 
weight of the carbon dioxide in the flask at the observed 
temperature and pressure; it may help you in seeing that 
this is so if you consider for a moment what would happen 
to the first weight of the flask if the air were all pumped out 
and a complete vacuum formed inside the flask before the 
carbon dioxide is let in. The calculations on the volume of 
the air in the flask apply equally well to carbon dioxide. 
What, then, would the volume be at S. T. P. of the weight 
of carbon dioxide you have just found? By dividing the 
weight of the carbon dioxide by its volume at S. T. P. you 
get the weight of 1 c.c. of carbon dioxide at S. T. P. How 
much is it ? You were told above that 1 c.c. of air at S. T. P. 
weighs 0.00129 gram. Divide the weight of 1 c.c. of car- 
bon dioxide at S. T. P. by the weight of 1 c.c. of air at 
S. T. P. and find the specific gravity of carbon dioxide re- 
ferred to air. 

Experiment 99. Combining number of magnesium. Have 

ready a 150 c.c. flask fitted with a good one-hole cork, or 
better, a rubber stopper, a 500 c.c. flask, graduate, ther- 
mometer, barometer, platform balance, set of smaller weights, 
pneumatic trough, and delivery tube; also some magnesium 
ribbon and hydrochloric acid. 

In the section on Dalton, under the Atomic Period, com- 



248 FIRST YEAR CHEMISTRY 

bining number was denned as the weight of one element that 
would unite chemically with unit weight of some other ele- 
ment, and it was stated that Dalton determined the com- 
bining numbers of several elements, referring them to hy- 
drogen, the lightest known element. More specifically the 
combining number of an element is the number of grams of 
that element which liberates, replaces or combines with one 
gram of hydrogen. The combining number is sometimes 
called the equivalent weight or simply the equivalent. It is 
often determined by allowing a metal to replace the hydro- 
gen of an acid. In such a case it might be called " replacing 
number," but combining number is the term ordinarily used. 

If we should take a large, weighed piece of magnesium 
and let an acid act on it till just one gram of hydrogen had 
been evolved, the loss in weight of the magnesium would 
show the amount of the latter which had entered into com- 
bination in the hydrogen's place, i.e., the amount of mag- 
nesium which is "equivalent" to one gram of hydrogen. 
To collect one gram of hydrogen would require an experi- 
ment on a large scale because this gas is so light. Hence, 
it is customary to start with a small, definite weight of the 
metal, measure the hydrogen evolved by it from an acid 
and then, by means of a proportion, calculate the com- 
bining number of the metal. 

Fit the 100 c.c. flask with a cork and a glass delivery tube 
leading to the pneumatic trough. Put in the flask 100 
c.c. of water and 10 c.c. of concentrated hydrochloric acid. 
Select a piece of clean, bright magnesium ribbon and weigh 
out a piece not less than 0.25 gram and not over 0.35 gram 
in weight. Weigh it accurately to hundredths of a gram 
and record this weight in the notebook. This weight of 
magnesium will produce a convenient amount of hydrogen 
to catch in the 500 c.c. flask. Invert the large flask full of 
water over the pneumatic trough. When all is ready, drop 



FIRST YEAR CHEMISTRY 



249 



the magnesium into the dilute acid in the small flask and 
quickly insert the stopper. Be sure to catch all the gas 
evolved. Let the reaction run to completeness. Keep all 
unnecessary heat away. Why ? Write the equation for the 
change that takes place. 

Note whether any water has been sucked back toward 
the generating flask. If any has, make proper allowance 
for it when measuring the gas caught. Note the tempera- 
ture of the air in the room; also read the barometer. Then 
make the level of the water inside and outside the large 
flask the same. Place the palm of the hand over the mouth 
of the flask and invert the flask. With the graduate add 
water to the brim and note how many c.c. of gas were col- 
lected in the flask. 

Correction for aqueous vapor. — Since the hydrogen was 
caught over water it contained considerable water vapor. 
The pressure inside the flask, tho it was equal to the atmos- 



Pressure of Aqueous Vapor 


Temp, in 

centigrade 

degrees 


Pressure 
in mm. of 
mercury 


Temp, in 

centigrade 

degrees 


Pressure 
in mm. of 
mercury 


Temp, in 

centigrade 

degrees 


Pressure 
in mm of 
mercury 


10 


9 


18 


15 


26 


25 


11 


10 


19 


16 


27 


27 


12 


10 


20 


17 


28 


28 


13 


11 


21 


18 


29 


30 


14 


12 


22 


20 


30 


32 


15 


13 


23 


21 


31 


33 


16 


13 


24 


22 


32 


35 


17 


14 


25 


23 


33 


37 



pheric pressure outside, was made up largely of the pressure 
due to the hydrogen, but to some extent of the pressure due 
to the water vapor. The true pressure of the hydrogen was, 



250 FIRST YEAR CHEMISTRY 

therefore, the atmospheric pressure diminished by the pres- 
sure of the aqueous vapor. Careful experimenters have 
found that the pressure exerted by water vapor increases 
with the temperature. Dalton showed that the pressure 
of water vapor could be measured by a mercury column in 
much the same way that the pressure of the air is measured. 
The accompanying table shows the correction that must be 
subtracted from the barometric reading at different tempera- 
tures when correcting the volume of a gas for aqueous vapor. 
Let us now find the weight of the hydrogen as follows: 
Find the true pressure on the hydrogen by subtracting from 
the barometric reading the proper correction from the above 
table. Using the Laws of Boyle and of Dalton, reduce the 
volume of hydrogen as measured to standard conditions. 
At standard conditions, i.e., at S. T. P., 1 c.c. of hydrogen 
weighs 0.00009 gram. Calculate the weight of the hydro- 
gen. Then form the following proportion: The weight of 
hydrogen evolved is to the weight of magnesium used as 
one gram of hydrogen is to the weight of magnesium needed 
to evolve one gram of hydrogen. The last term of this 
proportion is the combining number of magnesium. Sub- 
stitute in this proportion the weights used in your experi- 
ment and calculate the combining number of magnesium. 

Experiment 100. Determination of atomic weights. No 

apparatus is necessary for this experiment, because it is 
better for us to use the combining numbers obtained by 
more advanced workers than at present to determine these 
for ourselves. Furthermore, this experiment deals rather 
with the application of the combining numbers than with 
their determination. 

A little additional data on combining numbers may be 
of service at this point. Actual experiments have given 
the following results: (a) 1 gram of hydrogen unites di- 



FIRST YEAR CHEMISTRY 251 

rectly with 35.5 grams of chlorine to form hydrochloric acid 
gas; (b) 1 gram of hydrogen unites directly with 80 grams 
of bromine to form hydrobromic acid gas; (c) it requires ex- 
actly 23 grams of sodium to evolve one gram of hydrogen 
from water; (d) it requires exactly 39 grams of potassium to 
evolve 1 gram of hydrogen from water; (e) it requires ex- 
actly 12 grams of magnesium to evolve 1 gram of hydrogen 
from water; (f) 8 grams of oxygen unite directly with 1 
gram of hydrogen to form water; (g) 23 grams of sodium 
unite directly with 35.5 grams of chlorine to form sodium 
chloride; (h) 23 grams of sodium unite directly with 80 
grains of bromine to form sodium bromide; (i) 39 grams of 
potassium unite directly with 35.5 grams of chlorine to form 
potassium chloride; (j) 39 grams of potassium unite directly 
with SO grams of bromine to form potassium bromide; 
(k) 12 grams of magnesium unite directly with 35.5 grams 
of chlorine to form magnesium chloride; (1) 12 grams of 
magnesium unite directly with 80 grams of bromine to form 
magnesium bromide. All this may be condensed into the 
following list in which it will be easier to compare the figures: 

(a) weight of hydrogen : weight of chlorine = 1:35.5 

(b) weight of hy drogen : weight of bromine = 1 : 80 

(c) weight of hydrogen : weight of sodium = 1:23 

(d) weight of hydrogen : weight of potassium = 1 : 39 

(e) weight of hydrogen : weight of magnesium = 1:12 

(f) weight of hydrogen : weight of oxygen = 1:8 
(g) weight of sodium : weight of chlorine =23:35.5 

(h) weight of sodium : weight of bromine =23:80 

(i) weight of potassium : weight of chlorine =39:35.5 

(j) weight of potassium : weight of bromine =39:80 

(k) weight of magnesium : weight of chlorine =12:35.5 

(1) weight of magnesium : weight of bromine =12:80 

It will be noticed that in the first half of the above list the 



252 FIRST YEAR CHEMISTRY 

figures are the combining numbers of the different elements 
compared with hydrogen. It will be noticed also, that in 
the last half of the list these elements unite with each other 
in the ratio of their combining numbers. If these elements 
combine in each case atom to atom, then the combining 
numbers represent also the atomic weights of the elements. 
But we found, when considering the subject of valence, that 
in many cases an atom of one element will unite with two 
or more atoms of another element. This indicates that the 
atomic weight may often be a multiple of the combining 
number. Let us consider a specific case and see how we 
determine which multiple of the combining number to take 
as the true atomic weight.. 

In the preceding experiment we determined the combin- 
ing number of magnesium, i.e., the weight of magnesium 
that would unite with or rather replace unit weight of hy- 
drogen, and we found this to be about 12. If one atom of 
magnesium replaces one atom of hydrogen, then the atomic 
weight of magnesium will be 12, assuming that for hydrogen 
to be unity. But if one atom of magnesium replaces two 
atoms of hydrogen, then the atomic weight of magnesium 
must be twice 12 or 24. If all metals replaced hydrogen 
atom for atom, the atomic weights would always be identi- 
cal with the combining numbers. Since such is not the case, 
it follows that the atomic weights are sometimes the same 
as the combining numbers and sometimes multiples of the 
combining numbers. 

The specific heat of a metal helps us determine which 
multiple of the combining number to take as the atomic 
weight. By the specific heat of a substance we mean the 
amount of heat required to raise the temperature of that sub- 
stance one degree. As has been pointed out under the Atomic 
Period, Dulong and Petit, two French chemists, discov- 
ered that if the atomic weight of an element is multipled 



FIRST YEAR CHEMISTRY 



253 



by its specific heat the product usually comes out about 6.4; 
the accompanying table illustrates a few cases; it will be 
noted that the products vary from 6.0 to 6.8, but the aver- 
age is not far from 6.4. 



Illustration of Law of Dulong and Petit 


Element 


Specific heat 


Atomic weight 


Product of 

specific heat and 

atomic weight 


Lead 


0.0314 


207 


6.5 


Mercury 


0.0333 


200 


6.7 


Antimony 


0.0508 


120 


6.1 


Silver 


0.0560 


108 


6.0 


Zinc 


0.0955 


65.3 


6.2 


Iron 


0.1138 


56 


6.4 


Calcium 


0.1700 


40 


6.8 



Let us now see how the Law of Dulong and Petit helps 
us in determining the atomic weight of magnesium. If we 
divide the constant, 6.4, b} r the specific heat of magnesium, 
0.25, we get 26.0 as a quotient. Plainly, then, the combin- 
ing number, 12, of magnesium is not the atomic weight of 
this metal. The second multiple of the combining number, 
i.e., 24, is much nearer the. quotient, 26, than is either the 
combining number itself or its third multiple, 36. Therefore, 
we conclude that the second multiple of the combining num- 
ber represents the true atomic weight. In this particular 
case the atomic weight of magnesium has been determined 
in other ways, and the results confirm the value 24 as the 
true atomic weight. 

In the accompanying table will be found the specific heats 
and the combining numbers of several metals. The first 
case is that of magnesium just considered, and is worked out 
in full. From the data given in the table calculate the 
atomic weights of the metals. Copy the table with the 



254 



FIRST YEAR CHEMISTRY 



results of your calculations in the laboratory notebook un- 
der Experiment 100. 



Determination of Atomic Weights 


Element 


Specific 
heat 


6.4 divided 

by the 
specific heat 


Combining 
number 


Multiple of 
the combin- 
ing number 
nearest the 
quotient in 
the third 
column 


Atomic 
weight 












Magnesium 


0.25 


26.0 


12.0 


2 X 12 


24 


Potassium 


0.166 




39.0 






Sodium 


0.293 




23.0 






Tin 


0.055 




59.4 






Copper 


0.095 




31.6 






Calcium 


0.170 




19.9 






Aluminium 


0.214 




9.0 







Atomic weights. — Much careful work has been done on the 
determination of atomic weights, many chemists having made 
this their life work. On the following page will be found 
the atomic weights of the eighty-one elements now known. 
The figures given are only approximate, because these are 
the ones we shall use in our future calculations. For the 
exact atomic weights as revised from year to year the stu- 
dent is referred to the Report of the International Commit- 
tee on Atomic Weights, published annually in the January 
number of the Journal of the American Chemical Society. 

Determination of molecular weights. — We have just seen 
how atomic weights may be determined. Molecular weights 
are closely associated with atomic weights; in fact ; many 
atomic weights are determined by finding first the molecu- 
lar weight of some compound of the element under examina- 
tion, and then deriving the atomic weight of the element 
from the molecular weight of the compound. There are 






Table of Atomic Weights 


Aluminium 


. Al 


27 


Molybdenum 


. . Mo 


96 


Antimony . 


. Sb 


V?\) 


Neodymium 


. . Nd 


144 


Argon . . . 


. A 


40(?) 


Neon . . . 


. . Ne 


20 


Arsenic . . 


. As 


75 


Nickel . . . 


. . Ni 


58 


Barium . . 


. Ba 


137 


Nitrogen . . 


. . N 


14 


Bismuth . . 


. . Bi 


208 


Os?nium . . 


. . Os 


191 


Boron . . . 


. . B 


11 


Oxygen . . 


. . O 


16 


Bromine . . 


. Br 


80 


Palladium . 


. . Pd 


107 


Cadmium . 


. Cd 


112 


Phosphorus 


. . P 


31 


C cesium . . 


. Cs 


133 


Platinum 


. . Pt 


195 


Calcium . . 


. Ca 


40 


Potassium . 


. . K 


39 


Carbon . . 


. C 


12 


Praseodymium 


. Pr 


140 


Cerium . . 


. Ce 


140 


Radium . . 


. Ra 


225 


Chlorine . . 


. CI 


35.5 


Rhodium . . 


. . Rh 


103 


Chromium . 


. Cr 


52 


Rubidium . 


. . Rb 


86 


Cobalt . . . 


. Co 


59 


Ruthenium . 


. Ru 


102 


Columbium . 


. Cb 


94 


Samarium . 


. . Sm 


150 


Copper . . 


. Cu 


63.5 


Scandium . 


. . Sc 


44 


Dysprosium 


• Dy 


162 


Selenium . . 


. Se 


79 


Erbium . . 


. Er 


166 


Silicon . . . 


. Si 


28 


Europium . 


. Eu 


152 


Silver . . . 


- Ag 


108 


Fluorine . . 


. F 


19 


Sodium . . 


. Na 


23 


Gadolinium 


. Gd 


156 


Strontium . 


. Sr 


87 


Gallium . . 


. Ga 


70 


Sulphur . . 


. S 


32 


Germanium 


. Ge 


73 


Tantalum 


. Ta 


183 


Glucinum 


. Gl 


9 


Tellurium . 


. Te 


128 


Gold . . . 


. Au 


197 


Terbium . . 


. Tb 


160 


Helium . . 


. He 


4 


Thallium . . 


. . Tl 


204 


Hydrogen . 


. H 


1 


Thorium . . 


. Th 


233 


Indium . . 


. In 


115 


Thulium . . . 


. Tm 


171 


Iodine . . . 


. I 


127 


Tin 


. Sn 


119 


Iridium . . 


. Ir 


193 


Titanium . . 


. Ti 


48 


Iron .... 


. Fe 


56 


Tungsten . . 


. W 


184 


Krypton . . 


. Kr 


82 


Uranium . . 


. U 


239 


Lanthanum . 


. La 


139 


Vanadium . 


. V 


51 


Lead . . . 


. Pb 


207 


Xenon . . . 


. Xe 


128 


Lithium . . 


. Li 


7 


Ytterbium . 


. Yb 


173 


Lutecium 


. Lu 


174 


Yttrium . . 


. Yt 


89 


Magnesium . 


• Mg 


24 


Zinc . ... 


. Zn 


65. 5 


Manganese . 


. Mn 


55 


Zirconium . . 


. Zr 


91 


Mercury . . 


. Hg 


200 









256 FIRST YEAR CHEMISTRY 

two methods for determining molecular weights. The first 
is called the physical method, because it depends upon a 
physical property of the substance. The second is called the 
chemical method, because it depends upon some chemical 
change brought about in the substance under examination. 

Experiment 101. Determination of molecular weights by 
the physical method. No apparatus is necessary for this 
experiment, because it is better for us to take the specific 
gravities of the different gases as determined by more ad- 
vanced workers than for us to spend the large amount of time 
necessary in determining them. Furthermore, this experi- 
ment deals rather with the application of the specific gravi- 
ties than with their determination. 

This method depends upon the Suggestion of Avogadro, 
that equal volumes of all gases contain the same number 
of molecules. From this it follows that the ratio between 
the weights of single molecules of two different gases must 
be the same as the ratio between the weights of equal vol- 
umes of these two gases. When we have determined the 
specific gravity of a gas referred to hydrogen we may say, 
then, that the molecule of that gas is the same number of 
times heavier than the molecule of hydrogen. For instance, 
the specific gravity of carbon dioxide referred to hydrogen 
is 22; the molecule of carbon dioxide is, therefore, 22 
times heavier than the molecule of hydrogen. 

A careful study of the union between hydrogen and 
chlorine shows, however, that every molecule of hydrogen 
contains two atoms. That is why we write hydrogen gas 
H 2 and not simply H. Now we have already agreed to call 
the weight of the atom of hydrogen unity. It follows, 
therefore, that we must multiply the specific gravity of 
carbon dioxide by 2 in order to get the molecular weight 
of this gas; for, if the molecule of carbon dioxide is 22 times 



FIRST YEAR CHEMISTRY 257 

as heavy as the molecule of hydrogen, it must be twice 22 
(or 44) times as heavy as the atom of hydrogen; we say, 
then, that 44 is the molecular weight of carbon dioxide. 
This method is applicable only to those substances which 
are gases, or which can be converted into vapors. 

The figures given below indicate the specific gravities of 
the substances referred to hydrogen. Calculate and record 
in the laboratory notebook the molecular weights of the 
following different substances: 

Oxygen 16 Sulphurous Oxide 32 

Nitrogen 14 Carbon Monoxide 14 

Nitrous Oxide 22 Carbon Dioxide 22 

Nitric Oxide 15 Olefiant Gas 14 

Nitrogen Peroxide 23 Marsh Gas 8 

Chlorine 35.5 Acetylene 13 



Experiment 102. An experiment to determine molecular 
weights by the chemical method. Have ready a clean porce- 
lain crucible, pipe-stem triangle, tripod, Bunsen burner, 
horn-pan balance, set of smaller weights, and brass forceps; 
also some chemically pure, dry potassium chlorate. 

In the preceding experiment we found the molecular 
weight of oxygen to be 32. During our study of the lan- 
guage of chemistry we learned that the molecule of oxygen 
consists of two atoms of oxygen, hence, the atomic weight 
of oxygen must be 16. Let us now bring about a chemical 
change in which oxygen is involved, e.g., the heating of 
potassium chlorate. This will enable us to determine the 
molecular weights of potassium chlorate and of potassium 
chloride, and the atomic weight of potassium. 

Get the exact weight to centigrams of a clean, dry porce- 
lain crucible with its cover. Put in the crucible about 2 
grams of pure, dry potassium chlorate; put on the cover 



258 FIRST YEAR CHEMISTRY 

and get the exact weight to centigrams of the crucible, 
cover and contents. Heat the covered crucible with the 
Bunsen burner flame, first gently, and then with the full 
heat of the burner, taking care that the contents of the 
crucible do not foam up and boil over the sides. Lift the 
cover a little occasionally to see when the salt melts. After 
the chlorate has melted, continue heating to drive off the 
oxygen. Let the crucible and contents cool; then weigh it 
accurately to centigrams. Heat to constant weight, using 
the blast lamp if necessary. When the weight is constant, 
calculate the loss in weight due to the evolution of oxygen. 

Molecular weight of potassium chlorate. — To calculate the 
molecular weight of the chlorate, use the following propor- 
tion: The gram weight of the potassium chlorate is to the 
gram weight of the oxygen evolved as the molecular weight 
of the chlorate is to the molecular weight of the oxygen 
evolved. Let x equal the molecular weight of potassium 
chlorate. Every molecule of the chlorate yields three atoms 
of oxygen when heated; hence, our proportion becomes: 
The gram weight of the chlorate: the gram weight of the 
oxygen = x : 48. Calculate x. 

Molecular weight of potassium chloride. — Next make an- 
other proportion: The gram weight of the potassium chlo- 
rate is to the gram weight of the potassium chloride left in 
the crucible as the molecular weight of the chlorate deter- 
mined above is to the molecular weight of the potassium 
chloride left after the heating. One molecule of potassium 
chlorate produces one molecule of potassium chloride. From 
this calculate the molecular weight of potassium chloride. 

Atomic weight of potassium. — Finally, let us determine 
the atomic weight of potassium. We have just determined 
the molecular weight of potassium chloride. In the preced- 
ing experiment we calculated the molecular weight of chlo- 
rine and we know that it has two atoms to the molecule; 



FIRST YEAR CHEMISTRY 259 

from this we can determine the atomic weight of chlorine. 
The molecule of potassium chloride consists of one atom of 
potassium and one atom of chlorine. From this data cal- 
culate the atomic weight of potassium. 

Stoichiometry. — This name is given to calculating from 
the known weight of some factor or product the weights of all 
the other factors and products that enter into the reaction. The 
only helps needed are, (1) The correct equation for the re- 
action, (2) the atomic weights of the elements concerned, 
and (3) the ability to form an arithmetical proportion. 

The following example ma}^ make the above clear: Sup- 
pose a chemist needs just 100 grams of silver chloride and 
must prepare it from silver nitrate and calcium chloride. How 
man}j grams of silver nitrate must he start with to get just 
100 grams of silver chloride? How many grams of calcium 
chloride will be needed to bring about the change? How many 
grams of calcium nitrate will be formed as a side product? 
This problem should be solved as follows: First write the 
equation for the reaction that takes place: 

2AgN0 3 + CaCl 2 = 2AgCl + Ca(N0 3 ) 2 

Then write above the formula for silver chloride the amount 
wanted, e.g., 100 grams, and above the silver nitrate x, 
which is to represent the number of grams of silver nitrate 
needed, and above the calcium chloride y, which is to repre- 
sent the number of grams of the calcium salt needed, and 
above the calcium nitrate z, which is to represent the number 
of grams of calcium nitrate formed. Next, underneath each 
factor and product calculate its molecular weight. The 
molecular weight of a substance is always determined by 
adding together the atomic weights of all the atoms in the 
molecule; if more than one molecule of a substance is used, 
of course the molecular weights should be multiplied by 



260 FIRST YEAR CHEMISTRY 

the number of molecules used. The result will be some- 
thing like this: 



X IJ 


lOOg. z 


2AgN0 3 + CaCl 2 = 


- 2AgCl + Ca(N0 3 ) 2 


2(108 + 14+48) 40+2(35.5) 


2(108+35.5) 40+2(14+48) 


2(170) 40+71 


2(143.5) 40 + 124 


340 111 


287 164 



It is evident that the ratio between the silver nitrate start- 
ed with and the silver chloride obtained is always the same 
no matter what unit of measure is used, i.e., the ratio is 
always the same no matter whether it is expressed in grams 
or in molecular weights. In other words: The molecular 
weight of the silver chloride is to the molecular weight of 
the silver nitrate as the gram weight of the silver chloride 
is to the gram weight of the silver nitrate. By substitut- 
ing in this proportion we get the following numerical pro- 
portion : 

287 : 340 = 100 : x 
287x = 34000 
x = 118.46 g. 

The amount of calcium chloride needed may be found in 
a similar manner: 

287 : 111 = 100 : y 
287?/ = 11100 
V = 38.67 g. 

The amount of calcium nitrate formed may also be found 
in a similar manner: 

287 : 164 = 100 : z 
2S7z = 16400 
z = 57.14 g. 

A convenient check upon the accuracy of the arithmetical 



FIRST YEAR CHEMISTRY 261 

work, if the weights of all the substances are calculated, 
is to apply the Law of Conservation of Mass, and see if the 
sum of the weights of the factors does equal the sum of 
the weights of the products. In this problem 118.46 + 
38.67 = 100 + 57.14, or 157.13 = 157.14, which is within 
the limit of error caused by neglecting figures in the third 
decimal place. 

Rule of Stoichiometry. — From the above problem it is 
easy to formulate the following rule for working out prob- 
lems in stoichiometry: The molecular weight of the known 
is to the molecular weight of the unknown as the gram weight 
of the known is to the gram iceight of the unknown. The 
known substance is the one whose absolute weight in grams 
or pounds is given. Any member of the equation, either 
factor or product, may be considered the known substance; 
this, of course, is determined by the conditions of the ex- 
periment: then all the other members are for the time being 
the unknowns. 

Do the following problem on the left-hand page of the 
notebook, recording therein the proportions in full, as well 
as all the arithmetical work. 

Problem. — How many grams of nitric acid can be ob- 
tained from 500 grams of potassium nitrate by distillation 
with sulphuric acid? How many grams of sulphuric acid 
will be required? How many grams of potassium sulphate 
will be formed? Calculate all these values and check the 
accuracy of your work by applying the Law of Conserva- 
tion of Mass. H=l S=32 = 16 N = 14 K=39. Ans. 
311.8 g. HN0 3 , 242.5 g. H2SO4, 430.6 g. K 2 S0 4 . 

Application of stoichiometry to problems involving gas vol- 
umes. — If it is desired to find the volume of a gas that is 
evolved in a given case, the weight of that gas must first 
be determined by means of the rule of stoichiometry just 
explained above. From the weight of the gas it is easy to 



262 FIRST YEAR CHEMISTRY 

calculate its volume: (1) if the weight of one liter of it at 
the given conditions is known; or (2) if the weight of one 
liter of it at S. T. P. and the temperature and the pressure 
at the time of the experiment are given; or (3) we may have 
given the specific gravity of the gas referred to air or hydro- 
gen as a standard, and the weight of one liter of the stand- 
ard used. In the last two cases it will probably be neces- 
sary to apply the Laws of Boyle and of Dalton to the gas 
volume. 

Perhaps the following example might make this clearer: 
How many liters of carbon dioxide can be obtained from 500 
grams of calcium carbonate by means of hydrochloric acid 
when the gas is measured at 17° C. and 770 mm. pressure? 
The specific gravity of carbon dioxide is 1.53 referred to air, 
and one liter of air at S. T. P. weighs 1.293 grams. Ca=40 
C = 12 = 16. This problem should be solved as follows: 
First calculate the weight of carbon dioxide that is pro- 
duced. 



40 



500 g. 


X 


CaC0 3 + 2HC1 = CaCl 2 + H 2 


+ co 2 


-12 + 48 


12 + 32 


100 


44 


100 : 44 = 500 : x 




100 x = 22,000 




x = 220 grams. 





If one liter of air weighs 1.293 grams and the specific 
gravity of carbon dioxide is 1.53, then one liter of carbon 
dioxide must weigh 1.53 times 1.293 grams, or 1.98 grams 
at S. T. P. The 220 grams of carbon dioxide would occupy 
220 divided by 1.98 or 111.1 liters at S. T. P. and its vol- 
ume at 17°C. and 770 mm. pressure can be calculated by 
the customary proportions: 



FIRST YEAR CHEMISTRY 263 






V : V 


= P' : P 


V : V 


= 273+t :273+t' 


111.1 : V 


= 770 : 760 


109.6 : V 


= 273 : 273 + 17 


770V 


= 111.1X760 


109.6 : V 


= 273 : 290 


V 


111.1X760 


273V 


= .109.6X290 


740 


V 


109.6X290 


V 


= 109.6 


273 






V 


= 116.4 



The volume of carbon dioxide will, therefore, be 116.4 
liters at the conditions given in the problem. 

Do the following problem in the notebook, recording all 
the work. 

Problem. — How many liters of hydrogen gas can be ob- 
tained by treating 25 grams of zinc with sulphuric acid, 
the gas to be measured at 17°C. and 750 mm. pressure? 
One liter of hydrogen weighs 0.09 gram at 0°C. and 760 mm. 
pressure. Zn = 65.5 H = l. Ans. 9.081. 

Determination of the chemical formula from the percent- 
age composition of the compound. — It is customary to de- 
termine the correct molecular formula of a compound from 
its percentage composition and its molecular weight. The 
former is learned thru a quantitative analysis of the sub- 
stance, and the latter is ascertained in a number of ways, 
often from the vapor density. 

A recent college entrance examination paper contained 
the following problem: The percentage composition of al- 
cohol is carbon 52.18, hydrogen 13.04, oxygen 34.78. One 
liter of alcohol vapor at 136° and 760 mm. pressure weighs 
1.38 grams; one liter of hydrogen under the same conditions 
weighs 0.06 gram. Calculate the for mida of alcohol. C = 12 
H = 1 = 16. This problem should be solved as follows: 
If one liter of alcohol vapor weighs 1.38 grams and one liter 
of hydrogen weighs 0.06 gram, under the same conditions, 
the vapor density (or the specific gravity of the vapor) of 



264 



FIRST YEAR CHEMISTRY 



alcohol referred to hydrogen must equal 1.38 divided by 
0.06, or 23. The molecular weight of alcohol is twice its 
vapor density, or twice 23, i.e., 46. Of this molecular 
weight 52.18 per cent is carbon, i.e., 52.18 per cent of 
46, or 24, are carbon; as the carbon atom weighs 12 we see 
that there are two atoms of carbon in the molecule of alco- 
hol. Similarly for the other elements. The arithmetical 
work may be condensed as follows: 



Element 


Percent- 
age of 
element 
in com- 
pound 


Molecu- 
lar 
weight of 
com- 
pound 


Weight 
of ele- 
ment in 
com- 
pound 


Weight 
of ele- 
ment in 
com- 
pound 


No. of 
Atomic atoms of 
weight of element 
element in com- 
pound 


Carbon 

Hydrogen 

Oxygen 


0.5218 
0.1304 
0.3478 


X 46 = 
X 46 = 
X 46 = 


24.00 

5.99 

15.99 


24.00 

5.99 

15.99 


-f- 12 = 2 
-1 = 6 
-*■ 16 = 1 



Therefore, the formula for alcohol is C 2 H 6 0. This method 
is very important and is used extensively, because by its 
use the chemical formulae of all new compounds may be de- 
termined. 

Do the following problem in the notebook, recording all 
the work: 

Problem.— The specific gravity of a certain gas referred 
to hydrogen is 13.91. The analysis of this gas gives the 
following percentages: Carbon 85.62, and hydrogen 14.38. 
What is the formula of the gas ? C = 1 2 H = 1 . Ans. C 2 H 4 . 

Determination of the percentage composition of a com- 
pound from its formula. — The percentage of each element 
in a compound is usually determined by a quantitative 
analysis of the substance. Occasionally, however, it is de- 
sirable to calculate quickly the percentage composition of 
a compound when its molecular formula is known. For in- 
stance: Calculate the per cent of barium, of chlorine, and 



FIRST YEAR CHEMISTRY 265 

of water of crystallization in crystallized barium chloride, 
whose formula is BaCU2H 2 0. Ba = l'S7 Cl=35.5 H = 1 
= 16. This problem should be solved as follows: The 
molecular weight of crystallized barium chloride is 137 + 
(2 + 35.5) +2(2 + 16), or 137+71+36, or 244. Of this total, 
137 are barium, and the per cent of barium is easily ob- 
tained by dividing 137 by 244. Proceed similarly for chlo- 
rine and for the water of crystallization. It may make the 
solution clearer if the proportion form is used as follows: 

137 : 244 = x : 100 x = 0.5615 or 56.15% of Ba 

71 : 244 = y : 100 y = 0.2909 or 29.09% of CI 

36 : 244 = z : 100 z = 0.1475 or 14.75% of H 2 

By the same method the percentage composition of any 
compound may be calculated. 

Do the following problem in the notebook, recording all 
the work. 

Problem. — Calculate the per cent of water of crystalliza- 
tion in crystallized copper sulphate, CuS0 4 .5H 2 0. Cu=63.6 
S=32 = 16 H = l. Also calculate the per cent of S0 3 
in the same compound. Ans. 36.05% H 2 0. 32.05% S0 3 . 

Experimental work. — Let us now drop for a while the 
mathematical and strictly theoretical part of the course, 
and take up some more experimental and descriptive work; 
this work will be similar to that done during the first few 
months of the school year, but it will deal with the experi- 
ments in the light of the theoretical conceptions just stud- 
ied. Suppose we start with the halogens. 

The halogens are four in number, namely, chlorine, bro- 
mine, iodine, and fluorine. The name "halogen" means 
"salt forming"; this property is familiar in ordinary table 
salt, which was formed from the first halogen, — chlorine, — 
studied this year. The name "halogens" is applied to these 



266 FIRST YEAR CHEMISTRY 

four elements, because they all form salts which resemble 
sodium chloride. In all work with the halogens note the 
similarity between them. Note also any marked differ- 
ence in properties or behavior. 

Experiment 103. Bromine and its properties. Have 

ready a test tube rack with test tubes; also some bromine, 
ether, alcohol, sulphide of carbon, distilled water, litmus 
paper, and turmeric paper. 

Do all of this experiment in the hood. 

Examine bromine (which may be found in the hood) for 
its properties, particularly state, color, odor, ease of va- 
porizing, and effect on moist test papers. Use only the 
smallest possible amount of bromine on account of the cor- 
rosive action of the vapor of this substance on the mucous 
membrane of the nose. 

Test the solubility of bromine in various solvents as fol- 
lows: Fill a test tube with bromine vapors, but do not get 
any of the liquid into the test tube. Add 5 c.c. of distilled 
water, close the mouth of the test tube with the thumb, 
shake, and note the solubility of bromine in water. The 
resulting mixture is called bromine water, and is generally 
expressed Br 2 + Aq; Aq is an abbreviation for the Latin 
word aqua, meaning water, and is used to indicate an in- 
definite amount of water. There is no chemical union be- 
tween the bromine and the water in bromine water. In 
actual work bromine water is often used instead of liquid 
bromine. Heat the bromine water that you have made, 
and note any change in the color of the solution. What 
has happened? 

Next try the solubility of bromine in ether, and compare 
its solubility in this substance with its solubility in water 
as follows: Prepare some bromine water in a test tube from 
bromine vapors and 5 c.c. of water. Add 5 c.c. of ether 



FIRST YEAR CHEMISTRY 267 

slowly and without shaking. Ether does not mix with 
water, but separates by itself in a layer on top of the water. 
Note the color of each layer; then shake the test tube well 
to mix the contents. "When the mixed liquids have come 
to rest, note the color of the water in the bottom of the test 
tube and of the ether above it. In which solvent is bromine 
more soluble, ether or water? Make two sketches in the 
notebook of the test tube containing the mixed liquids. 
In the first sketch shade the liquid that contains the bro- 
mine before shaking, and in the second shade the liquid that 
contains the bromine after shaking. In each sketch indi- 
cate the composition of each layer of liquid. 

In a similar manner try the solubility of bromine in sul- 
phide of carbon, noting that this solvent is heavier than 
water. Note also the color of the sulphide of carbon solu- 
tion of bromine. Draw sketches as before. 

Lastly, try the solubility of bromine in alcohol. Since 
this solvent mixes with water in all proportions, it is best 
to add the alcohol directly to the bromine vapors in a test 
tube. 

Experiment 104. Hydrobromic acid and bromides. Have 
ready a test tube rack with test tubes and a Bunsen burner; 
also some hydrobromic acid, potassium bromide, and litmus 
paper. 

Hydrogen unites with bromine, but not as easily as with 
chlorine. The resulting hydrogen bromide, or hydrobromic 
acid, as it is sometimes called, is, like hydrogen chloride, a 
colorless gas and soluble in water; it is commonly used in 
aqueous solution, and it forms a series of compounds called 
bromides, which are similar in many respects to the chlo- 
rides. 

Examine some aqueous hydrobromic acid from the bottle 
on the shelf and note its properties. Heat 3 or 4 c.c. of the 



268 FIRST YEAR CHEMISTRY 

substance in a test tube, and compare it with hydrochloric 
acid in strength, odor, and ease of decomposing. Test 
with moist test papers both the liquid itself and the vapor 
arising from it when it is heated. Is there any change in 
the color of the solution? Explain the change. 

Examine a crystal of potassium bromide from the bottle 
on the shelf and get its chief properties. Compare this 
salt with the chlorides of sodium and of potassium, particu- 
larly in color, form, taste, and solubility in water. 

Experiment 105. Replacement of bromine in a bromide 
by chlorine. Have ready a test tube rack with test tubes; 
also potassium bromide, chlorine water, sulphide of carbon, 
and distilled water. 

Dissolve a small crystal of potassium bromide in about 
10 c.c. of distilled water in a test tube. Add from 5 to 10 
c.c. of " chlorine water" from the bottle on the shelf; this 
substance, which is sometimes written Cl 2 + Aq, is similar 
to bromine water and is made by passing an excess of chlo- 
rine gas thru cold water till this is saturated with the gas; 
chlorine water reacts similarly to gaseous chlorine with other 
substances and is more convenient to handle. 

Then add to the test tube 4 to 5 c.c. of sulphide of car- 
bon and shake the mixture. Allow the sulphide of carbon 
to settle; note its color, comparing this with the color of 
the sulphide of carbon that dissolved bromine in Experi- 
ment 103. What does this new color of the sulphide of 
carbon show in regard to the bromine? What pushed the 
bromine out of the bromide and took its place? The equa- 
tion for this change, neglecting the water of solution, may 
be written thus: 

2KBr + Cl 2 = Br 2 + 2KC1. 
Since water of solution was present it may be desirable to 



FIRST YEAR CHEMISTRY 269 

show this in the equation; water of solution is generally in- 
dicated by the abbreviation Aq; a set of brackets is often 
put around the Aq and the substance dissolved in it. Using 
Aq. the above equation may be written thus: 

[2KBr + Aq] + [Cl 2 +Aq] = [Br 2 +2KC1+Aq]. 

Since the symbol Aq represents an indefinite amount of 
water, it is allowable to write it only once on the right- 
hand side of the equation, tho it is used twice on the left- 
hand side. 

Experiment 106. Preparation of bromine. Have ready a 

dry test tube and a Bunsen burner; also some potassium 
bromide, black oxide of manganese, and some sulphuric 
acid. 

Mix a small amount of potassium bromide with about an 
equal quantity of black oxide of manganese in a test tube, 
and add a little concentrated sulphuric acid. If the reaction 
does not take place immediately, heat very gently and note 
the color of the bromine fumes arising from the mixture. 
If the color of the fumes is not prominent as you look thru 
the test tube in the ordinary manner, try looking straight 
down into the test tube; this little device often reveals the 
color when the ordinary method fails. 

The equation for the reaction is: 

2XaBr + Mn0 2 + 2H2S0 4 = Na 2 S0 4 + MnS0 4 + 2H 2 + Br 2 . 
This is the regular test for bromine in a bromide. Is manga- 
nese dioxide a catalytic agent in this reaction? 

Note on bromine and bromides. — Bromine itself does not occur free 
in nature, but several bromides — particularly those of sodium and of 
magnesium — do exist in sea water and in salt springs and salt deposits. 
Bromine is prepared from bromides by either of the two methods illus- 
trated by the two preceding experiments. The specific gravity of 
liquid bromine referred to water is 3.2, and of bromine vapors referred 



270 FIRST YEAR CHEMISTRY 

to air is 5.5; it boils at 60°C. Bromine is used in making compounds 
of bromine and in making dyes. 

Experiment 107. Iodine and its properties. Have ready 
a test tube rack with test tubes and a Bunsen burner; also 
some iodine, alcohol, ether, sulphide of carbon, distilled 
water, litmus paper, and turmeric paper. 

Examine a few crystals of iodine and get the properties 
of this substance. The color of iodine vapor is best studied 
by heating a small crystal of iodine in a dry test tube. Does 
iodine bleach test paper as did chlorine and bromine? 

Try the solubility of iodine in the following solvents: 
Water, ether, alcohol, and sulphide of carbon. This test is 
best made by taking equal portions of about 5 c.c. each of 
the four solvents separately in four clean test tubes and add- 
ing to each test tube one small crystal of iodine. Note 
whether or not the iodine dissolves; also note the intensity 
of color of the resulting solution. To any solvent that dis- 
solves the first crystal readily add some more iodine to get 
some idea as to the amount of iodine that it will dissolve. 
If cold water does not seem to dissolve iodine, heat the water 
a little. 

Experiment 108. Hydriodic acid and iodides. Have 

ready a test tube rack with test tubes, and a Bunsen burner; 
also some potassium iodide and distilled water. 

Hydrogen iodide or hydriodic acid, as it is sometimes called, 
is known, but it is an unstable substance, and cannot be 
easily made by the synthesis of hydrogen and iodine. It is 
similar to hydrogen chloride and hydrogen bromide, but it 
is much weaker. It forms iodides which are similar to bro- 
mides and chlorides. 

Examine a crystal of potassium iodide and get its chief 
properties. 









FIRST YEAR CHEMISTRY 271 

Experiment 109. Replacement of iodine in an iodide by 
chlorine. Have ready a test tube; also some potassium 
iodide, chlorine water, sulphide of carbon, and distilled water. 

Dissolve a small crystal of potassium iodide in about 10 
c.c. of water in a test tube. Add from 5 to 10 c.c. of chlo- 
rine water and 4 to 5 c.c. of sulphide of carbon. Shake the 
mixture and then let it settle. Note the color of the sul- 
phide of carbon. What does this show in regard to the 
iodine that was in the iodide of potassium? What pushed 
the iodine out of the iodide and took its place? Compare 
this with the action of chlorine on the bromide and write 
the equation for the replacement of iodine in an iodide by 
chlorine, neglecting the water of solution; then rewrite the 
equation showing the water of solution. 

Experiment 110. Replacement of iodine in an iodide by 
bromine. Have ready a test tube; also some potassium 
iodide, bromine water, and distilled water. 

According to Experiment 105, which had the stronger 
attraction for potassium, chlorine or bromine? According to 
Experiment 109, which had the stronger attraction for po- 
tassium, chlorine or iodine? Let us now see which has the 
stronger attraction for potassium, bromine or iodine. To 
a dilute solution of potassium iodide in a test tube add 
some bromine water, and then some sulphide of carbon. 
Shake and let settle. Judging from the color of the sul- 
phide of carbon what would you say has happened to the 
potassium iodide? Write the equation for the replacement 
of iodine in an iodide by bromine. 

Note on iodine and iodides. — Iodine itself does not occur free in na- 
ture, but several iodides — particularly those of sodium, of potassium, 
and of magnesium — do exist in sea water and in salt springs and salt 
deposits. Iodine may be prepared in the laboratory by a method 
.similar to the second method for preparing bromine; potassium iodide, 



272 FIRST YEAR CHEMISTRY 

black oxide of manganese, and concentrated sulphuric acid when heat- 
ed in a test tube evolve iodine vapors which condense on the sides 
of the test tube. The specific gravity of iodine is 4.9; it melts at 114°C, 
and boils near 184°C. Iodine is used in making dyes, and in making 
iodoform; the alcoholic solution of iodine is sold in trade under the 
name tincture of iodine and is used for medicinal purposes. 

Experiment 111. Calcium fluoride and its properties. 
Have ready two test tubes and a Bunsen burner; also some 
calcium fluoride. 

The fourth halogen, fluorine, is an extremely active ele- 
ment, so active that it cannot be handled in the elementary 
state. It unites with the greatest ease with hydrogen, form- 
ing hydrogen fluoride and this compound is even more ac- 
tive than the other halogen acids. For that reason we must 
start with calcium fluoride, a compound composed of cal- 
cium and fluorine. 

Examine some powdered calcium fluoride from the bottle 
on the shelf and note its chief properties, particularly its 
state, color, specific gravity, and solubility in water. 

Experiment 112. Hydrogen fluoride; its preparation, its 
properties, and its use in etching glass. Have ready a Bun- 
sen burner, tripod, gauze, filter paper, and two small beak- 
ers, of such sizes that the bottom of the larger may rest 
upon the rim of the smaller beaker without slipping down 
into it; - also some calcium fluoride, sulphuric acid, and paraf- 
fine. 

Warm the bottom of a small beaker and spread on its 
bottom, outside, a thin layer of paraffine. When this has 
cooled, trace some design or figure thru the paraffine by 
means of a pin or sharp file, being sure to cut thru the paraf- 
fine to the glass. 

In the other beaker to a little calcium fluoride add enough 
concentrated sulphuric acid to make a thick paste and heat 



FIRST YEAR CHEMISTRY 273 

it gently over the Bimsen burner. When the hydrogen 
fluoride begins to be evolved, note its color, state, action 
on moist test paper, and cautiously test its odor. Then 
set the paraffined beaker closely over the top of the other 
beaker and generate hydrogen fluoride for a few minutes. 
Be careful that the paraffine does not melt during the ex- 
periment. Finally remove the paraffined beaker, heat it 
a little, and remove the paraffine with filter paper. Ex- 
amine the surface of the glass to see if the hydrogen flu- 
oride has etched it. If the experiment was not a success, 
see if you can find out where the trouble lay and then try 
it again. Write the equation for the preparation of hydro- 
gen fluoride from calcium fluoride by means of sulphuric 
acid, assuming that the other product is calcium sulphate. 

Note on fluorine and fluorides. — As stated above, fluorine never oc- 
curs free in nature. The common fluorides found in nature are fluor- 
spar (calcium fluoride), and cryolite (a double fluoride of sodium and 
aluminium). Elementary fluorine has been made by electrolyzing 
hydrofluoric acid. The specific gravity of fluorine referred to air is 
1.3. There are no uses for fluorine. Hydrofluoric acid is used in 
etching glass. 

Experiment 113. Preparation of chlorine. Have ready 
a test tube rack with test tubes, and Bunsen burner; also 
some black oxide of manganese, hydrochloric acid, sodium 
chloride, sulphuric acid, bleaching powder, litmus paper, 
and turmeric paper. 

In the note on chlorine under Experiment 49 it was stated 
that there are three methods for making chlorine. Reread 
that note and then try the first two methods as follows: 

First Method. — Heat a little black oxide of manganese 
with some concentrated hydrochloric acid in a test tube. 
Smell cautiously of the chlorine evolved and verify its other 
properties. There are three products in the reaction, name- 
ly, manganese chloride, free chlorine, and water. Write 



274 



FIRST YEAR CHEMISTRY 



the equation, assuming that black oxide of manganese is 
manganese dioxide, Mn0 2 , but that the valence of manga- 
nese in manganese chloride is two. 

Modification of the First Method. — Mix in a dry test tube 
some sodium chloride and some manganese dioxide. Then 
add some concentrated sulphuric acid. Heat a little if 
necessary to start the evolution of the chlorine. This re- 
action runs as did the reaction for making bromine from 
sodium bromide by means of sulphuric acid and black oxide 
of manganese. Write the equation to represent this re- 
action. 

Note on the halogens. — Elements that resemble each other as much 
as do the four halogens are called a Natural Family of Elements. The 
accompanying table brings out the relation between the halogens 
clearly. Note particularly the gradation in the properties. 



Comparison of the Halogens. 


Properties 


Fluorine 
F 


Chlorine 
CI 


Bromine 
Br 


Iodine 
I 


Atomic weight 


19 


35.5 


80 


127 


Specific gravity 


1.3 (gas) 
1.14 (liquid) 


2.5 (gas) 
1.5 (liquid) 


5 (vapor) 
3 (liquid) 


8.7 (vapor) 
5.0 (solid) 


Boiling point 


— 187°C. 


— 33°C. 


60°C. 


184°C. 


State 


Gas 


Gas 


Liquid 


Solid 


Color 


Greenish 
yellow 


Greenish 
yellow 


Red brown 


Grayish 
black 


Action on 
Test paper 


Bleaches 
violently 


Bleaches 
strongly 


Bleaches 
slightly 


Does not 
bleach 


Union with 

hydrogen 

takes place 


In the 

dark at 

ordinary 

temperature 


On stand- 
ing in 
sunlight 


When 
heated 


At red 

heat, but 

incompletely 



Several other natural families exist and they will be referred to later. 



FIRST YEAR CHEMISTRY 275 

Second Method. — Put some bleaching powder in a test 
tube and add a little concentrated sulphuric acid. Note 
the chlorine evolved. Bleaching powder is a complicated 
compound of the probable composition, CaOCl2. The equa- 
tion for the above change is: 

CaOCl 2 + H2SO4 = CaS0 4 + Cl 2 + H 2 0. 

Hydrochloric acid may be substituted for sulphuric acid in 
the above reaction. Add a little concentrated hydrochloric 
acid to a little bleaching powder in a test tube and verify 
this statement. Then write the equation for the reaction 
between hydrochloric acid and bleaching powder. 

Experiment 114. Arsenic and its properties. Have ready 
a small test tube, Bunsen burner, and a magnifying glass; 
also some arsenic. 

Examine a small piece of arsenic and get its chief proper- 
ties. Compare it with the metals that you have studied. 
From its appearance would you consider it a metal? 

Put a small piece of arsenic in a small test tube and heat 
till the arsenic volatilizes and condenses on the sides of the 
tube. This deposit is called an "arsenic mirror." If your 
arsenic mirror is heavy, examine the inside of it with the 
magnifying glass, breaking the tube if necessary. 

Experiment 115. Oxidation of arsenic. Have ready a 
Bunsen burner, tripod, gauze, cover to a porcelain crucible, 
and magnifying glass; also some arsenic. 

Put a small piece of arsenic on the inverted cover of a 
porcelain crucible and heat it till the arsenic catches fire 
and burns. Note the color of the flame; also note the color 
and the odor of the fumes. Compare the arsenic oxide you 
made with some from the bottle on the shelf. Do not 
taste white arsenic, as it is a poison when taken internally. 



276 FIRST YEAR CHEMISTRY 

Oxide of arsenic is often called white arsenic. Its formula is 
As 2 3 , the valence of arsenic being 3. Write the equation 
for burning arsenic in air. 

Experiment 1 16. Reduction of arsenic oxide. Have ready 
a small test tube, porcelain mortar and pestle, and a Bun- 
sen burner; also some arsenic oxide and some charcoal. 

Mix equal portions of finely powdered charcoal and white 
oxide of arsenic, using only small portions of each. With 
a strip of folded filter paper clean the inside of the test tube 
from any adhering mixture. Warm gently over the Bunsen 
flame and note the arsenic mirror formed in the test tube. 
The carbon may be regarded as forming carbon dioxide, 
tho under some circumstances a little carbon monoxide 
may also be formed. Write the equation for the reduction 
of arsenic oxide by means of carbon, assuming that the prod- 
ucts are arsenic and carbon dioxide. Then write the equa- 
tion assuming that the products are arsenic and carbon mon- 
oxide. 

Experiment 117. Arsenic sulphide; its preparation and 
its properties. Have ready a test tube rack with test tubes, 
Bunsen burner, and a generator for making hydrogen sul- 
phide; also some arsenic oxide, hydrochloric acid, iron sul- 
phide, and sulphuric acid. 

First prepare a solution containing arsenic by dissolving 
a very little oxide of arsenic in a little concentrated hydro- 
chloric acid. Dilute the resulting solution of arsenic chlo- 
ride with about five time its volume of water. 

Into some of the clear arsenic chloride solution, pass a 
little hydrogen sulphide gas and note the precipitation of 
arsenic sulphide. This is the regular test for arsenic. Get 
the properties of arsenic sulphide, particularly its state, 
color, and solubility in water. Write the equation for the 



FIRST YEAR CHEMISTRY 277 

solution of arsenic oxide in hydrochloric acid; also write the 
equation for the preparation of arsenic sulphide from ar- 
senic chloride by means of hydrogen sulphide, assuming 
that the other product is hydrogen chloride. The formula 
for arsenic chloride is ASCI3, and for arsenic sulphide AS2S3. 

Note on arsenic and its compounds. — Arsenic is found free in nature, 
but it usually occurs combined with sulphur or a metal or with both. 
It may be prepared by roasting the ores in air and then reducing the 
resulting oxide of arsenic with carbon. The specific gravity of arsenic 
is 5.7; it sublimes without melting; at about 180°C. it burns in air. 
Its principal use is as an ingredient in lead shot. As stated above 
arsenic oxide is a vigorous poison; small doses of from 0.1 to 0.2 gram 
taken internally are usually fatal, tho workmen in arsenic factories 
often accidentally swallow larger portions without fatal results. Freshly 
precipitated iron hydrate is a good antidote; arsenic oxide is used in 
making rat and fly poisons, as a preservative for skins, and in medi- 
cines. Scheele's green is a fine, light green powder containing copper, 
arsenic, and oxygen, and called copper arsenite; it was formerly widely 
used in green paint and green wall paper. Paris green is also a light 
green powder, but more complicated than Scheele's green; it may be 
considered a mixture of copper arsenite and copper acetate; it is used 
as a bug and insect poison. 

Experiment 118. Antimony and its properties. Have 
ready a small test tube and a Bunsen burner; also some anti- 
mony, — both lump and powder. 

Examine both forms of antimony and get its chief proper- 
ties. Compare antimony in all respects with arsenic. Heat 
a little antimony in a small test tube and see if you can get 
a deposit similar to the arsenic mirror. 

Experiment 119. Oxidation of antimony. Have ready a 
Bunsen burner, tripod, gauze, cover of a porcelain crucible, 
and magnifying glass; also some antimony. 

Try to oxidize antimony on the cover of a porcelain cruci- 
ble, as you oxidized arsenic. Does the antimony burn? 



278 FIRST YEAR CHEMISTRY 

If it does, note the color of the flame; also note the color 
and the odor of the fumes. 

In order to oxidize antimony more readily heat a small 
lump of the metal in a very small blast lamp flame. Note 
the fumes of antimony oxide that arise. Suddenly remove 
the flame and watch the globule of melted antimony as it 
becomes covered with oxide. Examine the resulting sub- 
stance with the magnifying glass. 

Hold a small lump of antimony in the blast lamp flame, 
and when it melts let it fall thru a distance of at least 25 or 
30 cm. upon a large sheet of paper. Note the peculiar break- 
ing apart of the antimony globule into small globules; note 
also the tracks left by the small globules on the paper. 

Write the equation for the oxidation of antimony, remem- 
bering that this is a trivalent element. 

Experiment 120. Antimony sulphide ; its preparation and 
its properties. Have ready a test tube rack with test tubes, 
Bunsen burner, and a generator for making hydrogen sulphide; 
also some tartar emetic, iron sulphide, and sulphuric acid. 

Antimony does not react readily with acids; hence, solu- 
tions containing antimony are not easy to make. The most 
common soluble antimony compound is tartar emetic, a 
complicated compound containing potassium, antimony, 
and the tartaric acid group; its formula is KSbO.C 4 H 4 6 ; 
this salt reacts with hydrogen sulphide as you might expect 
antimony chloride to react if it were convenient to get the 
latter compound. Dissolve a very little tartar emetic from 
the bottle on the shelf, in half a test tube of water and pass 
in some hydrogen sulphide. Note the precipitation of anti- 
mony sulphide. This is the regular test for antimony. Get 
the chief properties of antimony sulphide. 

Note on antimony. — The chief ore of antimony found in nature is 
stibnite, antimony sulphide. The element is prepared by roasting the 






FIRST YEAR CHEMISTRY 279 

stibnite in air and reducing the resulting oxide with carbon. The 
specific gravity of antimony is 6.7; it melts at 430°C, and boils at 
1500°C. Unlike most metals, it expands on cooling, and takes a sharp 
impression of the mold; it is, therefore, used in type metal. 

Experiment 121. Lead and its properties. Have ready a 
Bunsen burner; also some lead, — both mossy and sheet. 

Examine some lead in the two forms, sheet and mossy, 
and get its chief properties, comparing it with other metals 
that you have studied. 

Note on lead. — Several compounds of lead occur as minerals, but 
the most common one is galena, lead sulphide. To obtain metallic 
lead the galena is roasted, first in contact with air and then out of 
contact with air; dining the first heating part of the lead sul- 
phide is oxidized to lead oxide and to lead sulphate, and during the 
second heating these compounds react with the remaining lead sul- 
phide, yielding metallic lead. The specific gravity of lead is 11.4; it 
melts at 326°C. and boils at 1700°C. The uses of lead are numerous; 
lead pipe and lead sheet are used widely for building purposes, in 
electrical work, and in making sulphuric acid; many alloys, such as 
type metal and solder, contain lead. 

Experiment 122. Oxidation of lead. Have ready a small 
Hessian crucible, ring stand and small ring, iron rod, and 
blast lamp; also some mossy lead. 

Fill a small Hessian crucible about half full of mossy 
lead, and heat it over the blast lamp, stirring occasionally 
with an iron rod to aid the oxidation. Continue the heat- 
ing till you get a considerable quantity of the oxide. Note 
that the lead forms several oxides of different colors. Does 
the lead burn with a flame as did zinc under similar condi- 
tions? What is the general color of the resulting powdered 
lead oxide that you get? Examine the two oxides of lead 
from the bottles on the shelf. Write the equation for the 
oxidation of lead, assuming that you got lead monoxide 
only. 



280 FIRST YEAR CHEMISTRY 

Note on the oxides of lead. — One of the two most common oxides 
of lead is lead monoxide. PbO, commonly called litharge or massicot; 
the other is lead tetroxide with the formula PD3O4, and is commonly 
called red lead or minium. There are three other oxides of lead known, 
PD2O, PD2O3, and Pb02, but they are not so common. Litharge is 
used in making other compounds of lead and in making glass; red 
lead is used as a paint and in plumbing. 

Experiment 123. Reduction of lead oxide. Have ready 
a porcelain mortar and pestle, a brick, and the blast lamp; 
also some litharge and some lampblack. 

Grind together in the mortar equal portions of about 5 c.c. 
each of litharge (lead monoxide) and lampblack (carbon). 
Empty the mixture out on a brick and play the blast lamp 
flame directly upon the mass. [After about 5 minutes heating, 
let the mass cool. What has happened to the lead mon- 
oxide ? Write the equation for the reduction, assuming that 
carbon dioxide was the other product. 

Experiment 124. Action of acids on lead. Have ready a 
test tube rack with test tubes, graduate, three small bottles 
or flasks, labels, and a Bunsen burner; also some mossy 
lead, hydrochloric acid, sulphuric acid, nitric acid, and dis- 
tilled water. 

First, prepare for use in this and some of the following 
experiments, a dilute solution of hydrochloric acid of a 
strength of 1 to 5 by adding 10 c.c. of concentrated hydro- 
chloric acid to 50 c.c. of distilled water; keep it in a labelled 
bottle or flask. Then prepare a dilute solution of sulphuric 
acid of a strength of 1 to 5, and a nitric acid solution of 
1 to 5 strength. 

Read the rest of the experiment thru before doing any more 
work. 

Put a few pieces of mossy lead in a test tube and add a 
little dilute hydrochloric acid of 1 to 5 strength. Is there 



FIRST YEAR CHEMISTRY 



281 



any action in the cold, i.e., at ordinary temperature? Heat 
the test tube a little. Is there any action in the hot? In 
another test tube try the action of lead with concentrated 
hydrochloric acid, first, in the cold and then in the hot. 

Repeat all this work, using sulphuric acid, i.e., try the 
action on lead of dilute sulphuric acid, first, cold, then hot, 
then the action of concentrated sulphuric acid, first, cold, 
then hot. 

Finally repeat all this work, using nitric acid. 

Add considerable distilled water to any white precipi- 
tate that may have formed in any of the above cases and 
see if it is soluble in water. Arrange the results in the 
form of a table as shown in the accompanying form. State, 
in your record, whether the action is vigorous, weak, or 



Action of Acids on Lead 


Conditions 


HC1 


H 2 S0 4 


HNO3 


Cold dilute 








Hot dilute 








Cold concentrated 








Hot concentrated 







wanting; it is always well in case there seems to be an ac- 
tion, to remove the test tube from the flame and watch to 
see if the action continues; sometimes, particularly in the 
case of hydrochloric acid, a seemingly vigorous action may 
be only the hydrochloric acid gas escaping from the hot 
liquid; in that case the bubbles cease quickly when the test 
tube is removed from the flame. If a gas is evolved, record 



282 FIRST YEAR CHEMISTRY 

that fact; also try to determine the gas. Record the color 
of the resulting solution. If any sediment is formed in the 
test tube, record that fact; record also the solubility in 
water of that sediment. If you are in doubt whether any of 
the metal has gone into solution, evaporate a few drops of 
the liquid to dryness in a clean test tube. 

From the results recorded in your table in the notebook, 
try to write the equations for the action of the different 
acids on lead. 

Experiment 125. Preparation of lead chloride by meta- 
thesis. Have ready a test tube rack with test tubes, beaker, 
tripod, gauze, graduate, and Bunsen burner; also some lead 
nitrate, hydrochloric acid, and distilled water. 

First prepare for use in this experiment, and those im- 
mediately following it, a dilute lead nitrate solution by dis- 
solving about 25 grams of lead nitrate in about 100 c.c. of 
distilled water, heating, if necessary, to aid the dissolving. 

Take about 20 c.c. of lead nitrate solution in a test tube, 
making sure the solution is cool, and then add some dilute 
hydrochloric acid. Note the precipitation of lead chlo- 
ride, and get its properties. Write the equation for the 
metathesis between lead nitrate and hydrochloric acid, as- 
suming that the other product is nitric acid. 

Let the precipitated lead chloride settle; then decant the 
clear, supernatant liquid and throw it away. Add about 
half a test tube of water to the lead chloride and heat till 
the precipitate has all dissolved. If the precipitate does 
not dissolve in the half a test tube of water on boiling, 
keep on adding a little more water and heating until the 
precipitate does dissolve. Allow the solution to cool and 
compare the crystallized lead chloride that separates from 
the solution with the amorphous form produced by precipi- 
tation. 



FIRST YEAR CHEMISTRY 283 

Experiment 126. Preparation of lead sulphate by meta- 
thesis. Have ready a test tube rack with test tubes and a 
Bunsen burner; also some sulphuric acid, distilled water, 
and some of the lead nitrate solution prepared in the pre- 
ceding experiment. 

Take about 20 c.c. of the prepared lead nitrate solution 
in a test tube. This time add some dilute sulphuric acid 
and note the precipitation of lead sulphate. The other 
product is nitric acid; write the equation for the metathesis. 
Get the properties of lead suphate. 

Experiment 127. The action of heat on lead nitrate. Have 

ready a test tube and a Bunsen burner; also some solid lead 
nitrate. 

Heat a few lumps of dry lead nitrate in a dry test tube 
and record all phenomena. The snapping of the crystalline 
solid is called decrepitation. Decrepitation is the violent 
snapping apart of a solid when it is heated, this snapping be- 
ing due to the explosive escape of minute particles of water 
mechanically inclosed in the solid. Judging from the color 
of the gas evolved, what gas is set free during the heating? 
Judging from the color of the residue in the test tube after 
cooling, what solid results from the heating? Write the 
equation for the effect of heat on lead nitrate, assuming 
that in addition to the two products just noted there is a 
third product, oxygen. 

Note on heating nitrates in general. — Roughly speaking, all common 
nitrates, except those of sodium, potassium, and ammonium, break 
down under the influence of heat into the oxide of the metal, nitrogen 
dioxide, and free oxygen. 

Experiment 128. Replacement of lead by zinc. Have 
ready a test tube; also some sheet zinc and some lead nitrate 
)lution. 



284 FIRST YEAR CHEMISTRY 

Take about 20 c.c. of the lead nitrate solution in a test 
tube. Into the clear solution insert a narrow strip of sheet 
zinc and allow it to stand for some time. Note the deposit 
of powdered lead on the surface of the zinc. This deposit 
is called the lead tree. Write the equation for the 
change. 

Experiment 129. Preparation of lead carbonate. Have 
ready a test tube; also some of your prepared lead nitrate 
solution and some ammonium carbonate solution from the 
bottle on the shelf. 

To about 20 c.c. of lead nitrate solution in a test tube 
add a little ammonium carbonate solution, (NH 4 ) 2 C0 3 . 
Note the precipitation of lead carbonate, and get its proper- 
ties. Write the equation for the metathesis between lead 
nitrate and ammonium carbonate, assuming that the other 
product is ammonium nitrate, NH4NO3. 

Note on lead carbonate. — Ammonium carbonate is the only soluble 
carbonate that will precipitate lead carbonate (PbCOs) from a lead 
solution. Sodium carbonate and potassium carbonate precipitate 
from a lead solution a basic carbonate, the composition of which varies 
with the temperature and the concentration of the solution. The most 
important of these basic carbonates is 2PbC03.Pb(OH) 2 ; this compound 
is the ' white lead used by painters and it may be made by allowing 
sheet lead, vinegar, and tan bark to stand in contact with each other 
for several months. 

Experiment 130. Preparation of lead chromate. Have 
ready a test tube; also some of your prepared lead nitrate 
solution and some potassium chromate solution from the 
bottle on the shelf. 

To about 20 c.c. of lead nitrate solution in a test tube 
add a little potassium chromate solution (K 2 Cr0 4 ). Note 
the precipitation of lead chromate and get its properties. 
Write the equation for the metathesis between lead nitrate 



FIRST YEAR CHEMISTRY 285 

and potassium chromate, assuming that the other product 
is potassium nitrate. 

Note on lead chromate and chromium compounds. — Lead chromate 
is called chrome yellow, and is used extensively as a yellow paint. 
Chromium is a gray, lustrous metallic element. It is different from 
most elements already studied, in that it unites with acid radicals to 
form a series of salts, e.g., chromium chloride, chromium sulphate, 
etc.: it also combines with oxygen to form a chromium acid radical 
that unites with metals to form chromates. All chromium compounds 
are colored. 

Experiment 131. Tin and its properties. Have ready a 
Bunsen burner and blast lamp; also some mossy tin, stick 
tin. block tin, tin foil, and tin plate. 

Examine the various forms of tin, such as mossy tin, 
stick tin, block tin, tin foil, and tin plate. Get the chief 
properties of this metal, particularly its color, luster, per- 
manency of luster, ease of melting, and hardness; also the 
"cry," of tin, i.e., the grating or crackling sound when a 
piece of mossy tin is bitten by the teeth or when a stick of 
tin is bent; this cry is supposed to be caused by the scrap- 
ing of the tin crystals over each other. 

Note on tin. — Tin seldom occurs free in nature. The principal ore 
is tin stone (tin dioxide), and from this ore metallic tin may be ob- 
tained by reduction with carbon. The specific gravity of tin is 7.3; 
it melts at 230°C, and boils at about 1700°C. On account of its sta- 
bility in air it is used in tin plating, i.e., dipping sheet iron into melted 
tin. Many alloys .such as solder, pewter, and "white metal" contain 
tin; these will be considered more fully in a later experiment on alloys. 

Experiment 132. Oxidation of tin. Have ready a brick, 
and the blast lamp; also some mossy tin. 

Put a piece of mossy tin on the brick and heat it with 
the blast lamp. Compare the ease of oxidizing tin with 
the ease of oxidizing lead. If you get any tin oxide note 
its properties and write the equation for forming it. 



286 FIRST YEAR CHEMISTRY 

Experiment 133. Action of acids on tin. Have ready a 
test tube rack with test tubes and a Bunsen burner; also 
some mossy tin, hydrochloric acid, sulphuric acid, and 
nitric acid, both the concentrated acids and the dilute acids 
you prepared some time ago. 

Treat small pieces of mossy tin in separate test tubes 
with the three acids, hydrochloric, sulphuric, and nitric, — 
dilute, concentrated, cold, and hot. Arrange the results in 
the form of a table as in the work on lead. In writing the 
equations consider tin as a bivalent metal. 

Experiment 134. Aluminium and its properties. Have 

ready a Bunsen burner and the blast lamp; also aluminium 
in its various forms, sheet, wire, foil, powder, and ingot. 

Examine the different forms of this metal, sheet, wire, 
foil, powder and ingot, and note its chief properties. Com- 
pare this metal with the other metals that you have studied. 

Note on aluminium. — This metal never occurs free in nature, but its 
compounds are numerous, abundant, and widely distributed. It is 
commonly found as a silicate of aluminium and of other metals in 
such rocks and minerals as clay, slate, feldspar, mica, and garnet; 
the oxide and hydrate are also found in considerable quantity, as is 
the cryolite mentioned in the note under fluorine. Metallic aluminium 
is obtained at present almost entirely by the electrolysis of aluminium 
hydrate dissolved in cryolite. The specific gravity of aluminium is 
2.6; it melts at 660°C, and boils at 1470-1700°C. The uses of alu- 
minium are numerous; it is used wherever lightness of weight is needed 
coupled with metallic strength; the powder is used in paints that are 
to be exposed to high heat, such as on radiators and steam pipes; the 
wire has recently been introduced as a conductor of electricity; large 
quantities of aluminium are used to reduce metallic oxides. 

Experiment 135. Oxidation of aluminium. Heat some 
aluminium sheet or wire, first in the Bunsen flame and then 
in the blast lamp flame. Does it oxidize easily? Alumin- 
ium oxide is known, and its formula is A1 2 3 . It is made 



FIRST YEAR CHEMISTRY 287 

more easily by roundabout methods from aluminium com- 
pounds, particularly the hydrate. Examine both varieties 
of aluminium oxide from the bottles on the shelf and get 
its properties. 

Note on aluminium oxide. — Aluminium oxide is often called alumina ; 
native crystallized varieties of alumina are called ruby, sapphire, and 
corundum ; impure and massive corundum is called emery. The last 
named variety is used for grinding and polishing; the others are used 



Experiment 136. Action of acids on aluminium. Have 
ready a test tube rack with test tubes, and a Bunsen burner; 
also some aluminium wire, hydrochloric acid, sulphuric acid, 
and nitric acid, both the concentrated acids and the dilute 
acids you prepared some time ago. 

Treat small pieces of aluminium in separate test tubes 
with the three acids, hydrochloric, sulphuric, and nitric, 
dilute and concentrated, cold and hot. If the action of 
aluminium with any one of the dilute acids is too violent 
for convenience, dilute the acid still more with water. Test 
the gas evolved. Arrange the results in the form of a 
table as in the work on lead. Write the equations for the 
action of acids on aluminium. Evaporate the resulting 
solutions to dryness and examine the residues. Compare 
the aluminium salts that you have made with as many 
aluminium salts as you can find on the shelf. 

Experiment 137. Alum and its properties. Have read}" 
two medium beakers, tripod, gauze, Bunsen burner, horn- 
pan balance, set of smaller weights, and brass forceps; 
also some potassium sulphate, aluminium sulphate, and alum. 

Weigh out exactly 10 grams of potassium sulphate and 
20 grams on aluminium sulphate from the bottles on the 
shelf. Dissolve eadi in as little water as will dissolve it, 



288 FIRST YEAR CHEMISTRY 

heating to aid the solution. Mix the two solutions and set 
away to crystallize. The resulting salt, alum, often crys- 
tallizes better than any of the other salts that we have 
made this year. The crystals that you get will probably 
show the crystalline form better than does the alum on the 
shelf, for this consists of broken fragments of larger crystals. 

Note on alum. — Alum is a double sulphate of aluminium and potas- 
sium, and crystallizes with 24 molecules of water of crystallization' 
its formula is, therefore, K^SO^A^SO^^r^O. This is often written 
K 2 A1 2 (S04)4.24H20. Alum is a very common compound of alumini- 
um. It is used in dyeing and printing cloth, in tanning, in paper 
making, in making other compounds of aluminium, in rendering wood 
and cloth fireproof, and as a medicine. 

Note on alums in general. — The alum just studied is often called 
potash alum, and it is typical of a large number of double salts of sim- 
ilar composition. They are all called alums, even tho several contain 
no aluminium whatever. They all crystallize in cubes or octahedra. 
Some of the more important alums are: 

K 2 A1 2 (S0 4 ) 4 .24H 2 potash alum 

Na 2 Al 2 (S0 4 )4.24H 2 soda alum 

(NH 4 ) 2 Al2(S0 4 ) 4 .24H 2 ammonium alum 

K 2 Cr 2 (S0 4 ) 4 .24H 2 potash chrome alum 

(NH 4 ) 2 Fe 2 (S0 4 ) 4 .24H 2 ammonium iron alum 

Experiment 138. Aluminium hydrate ; its preparation and 
its properties. Have ready a test tube rack with test tubes, 
and a Bunsen burner; also some alum, sodium hydrate, and 
distilled water. 

Dissolve a small lump of alum in half a test tube of 
water and add to it a few drops of dilute sodium hydrate 
solution. Note the white, voluminous precipitate of alu- 
minium hydrate, A1(0H) 3 . Write the equation, assuming 
that potassium sulphate and sodium sulphate are the other 
products. 

Note on insoluble hydrates. — This compound, Al(OH) 3 , introduces 
us to a large class of insoluble hydrates, i. e., the hydrates of the heavy 



FIRST YEAR CHEMISTRY 289 

metals. These are generally colored compounds, and may be precipi- 
tated from solutions of the metals by dilute sodium hydrate solution. 

Experiment 139. Silver; its properties and its oxidation. 
Have ready a Bunsen burner; also a piece of sheet silver. 

Examine a piece of sheet silver and note its chief proper- 
ties, comparing it with other metals that you have studied. 
Try to oxidize silver as you have oxidized other metals. 

Note on silver. — Silver occurs free in nature; its most important ore 
is argentite (silver sulphide); it also occurs in large quantities associ- 
ated with sulphur, arsenic, antimony, copper, and other metals. The 
extraction is complicated, and varies with the percentage of silver in 
the ore as well as with the constitution of the ore. Roughly speak- 
ing, the silver ore is roasted to remove the impurities. The resulting 
alloy of silver, lead, and gold is treated with zinc which alloys with 
the silver, and from this silver-zinc alloy the zinc is distilled, leaving 
the silver. Some ores of silver are treated successively with sodium 
chloride, iron, and mercury, and from the resulting silver-mercury 
amalgam, the silver is obtained by distilling off the mercury. The 
specific gravity of silver is 10.5; it melts at 960°C, and boils at a white 
heat. Silver is used in making silver coins and in silver plating. 

Experiment 140. Action of acids on silver. Have ready 
a test tube rack with test tubes and a Bunsen burner; also 
some sheet silver, hydrochloric acid, sulphuric acid, and 
nitric acid, both the concentrated acids and the dilute acids 
you prepared some time ago. 

Treat small pieces of sheet silver in separate test tubes 
with the three acids, hydrochloric, sulphuric and nitric, 
dilute and concentrated, first at ordinary temperature, 
and then with continued boiling. Arrange the results in 
the form of a table as in the work on lead. Get the proper- 
ties of any silver salts that you succeed in making. Write 
the equations for those cases where you find any action. 

Note on silver nitrate. — Silver nitrate is the most common soluble salt of 
silver. It is called lunar caustic. It is used in making other salts of silver. 



290 FIRST YEAR CHEMISTRY 

Experiment 141. Halogen salts of silver. Have ready a 
test tube rack with test tubes; also some silver nitrate solu- 
tion from the reagent bottle on the shelf, distilled water, 
some solid calcium chloride, potassium bromide, and potas- 
sium iodide. 

Put equal portions of about 5 c.c. each of silver nitrate 
solution in three test tubes. In three other test tubes pre- 
pare (1) a solution of calcium chloride, (2) a solution of 
potassium bromide, and (3) a solution of potassium iodide, 
using very small portions of each salt and dissolving them 
in distilled water. 

To one portion of silver nitrate add some calcium chloride 
solution. Note the properties of the precipitated silver chlo- 
ride. Shake the test tube to see if the precipitate clots. 
Let it stand exposed to light for some time to see if its color 
changes. Write the equation for the metathesis, assuming 
that calcium nitrate stays in solution. 

To another portion of silver nitrate solution add some 
potassium bromide solution and study the precipitated silver 
bromide in the same manner that you studied silver chloride. 
Write the equation for the metathesis, assuming that potas- 
sium nitrate stays in solution. 

To the last portion of silver nitrate solution add some 
potassium iodide solution and examine the precipitated 
silver iodide. Write the equation for the metathesis, as- 
suming that potassium nitrate stays in solution. 

Note on photography. — The halogen salts of silver are used exten- 
sively in photography; this industry is based on the fact that silver 
bromide and iodide change when exposed to light. These salts are 
held in a thin layer of gelatine on a glass plate or flexible gelatine 
film ; when this is exposed to light the light-struck parts are changed 
in proportion to the intensity of the light. The plate is then devel- 
oped, i.e., treated with a reducing agent that changes the light-struck 
parts to metallic silver. The unchanged silver salts are then dissolved 
out by "hypo," a solution of sodium hyposulphite, thus fixing the 



FIRST YEAR CHEMISTRY 



291 



image and leaving the negative, which shows the light places of the 
object as dark on the plate and vice versa. See Figs. 79 and 80. Print- 
ing is practically a repetition of the above; paper coated with silver 
salts is covered with the negative and exposed to light; the negative 
obstructs the light in proportion to the thickness of the silver deposit 




Fig. 79. A photographic positive (print). 



P- 






a 




Fig. 80. A photographic negative (plate). 

so the photograph has the same shading as the original object. The 
print, like the plate, must be treated with developer and hypo. Some- 
times the color of the print is improved by toning, i.e., by putting 
the print in a solution of gold or platinum. 



Experiment 142. Silver sulphide ; its preparation and its 
properties. Have ready a test tube rack with test tubes, 



292 FIRST YEAR CHEMISTRY 

Bunsen burner, and a generator for making hydrogen sul- 
phide; also some metallic silver, silver nitrate solution, iron 
sulphide, sulphuric acid, and nitric acid. 

Put some silver nitrate solution in a test tube and pass 
into it a little sulphide of hydrogen from the generator. 
Get the properties of the precipitated silver sulphide, par- 
ticularly its state, color, solubility in water, and solubility 
in hot, concentrated nitric acid. 

Write the equation for the precipitation of silver sulphide, 
assuming that the nitric acid is the other product. Write 
the equation for the solution of silver sulphide in concen- 
trated nitric acid; there are four products, namely, silver 
nitrate, nitric oxide, water, and free sulphur; the last prod- 
uct separates as a yellowish insoluble residue. 

If you have some metallic silver left, hold a piece of it in 
the stream of hydrogen sulphide. Does any deposit form? 
What is this deposit? Hold a dime, which is largely silver, 
in this stream of hydrogen sulphide. Do you get the same 
deposit here? The so-called oxidized silver ornaments of 
trade are not silver bearing a deposit of silver oxide, but 
silver bearing a deposit of silver sulphide. 

Experiment 143. Replacement of silver by copper. Have 

ready a test tube rack with test tubes; also some silver ni- 
trate solution, a narrow strip of copper, nitric acid, and dis- 
tilled water. 

Take about half a test tube of silver nitrate solution, 
and put into it a strip of freshly cleaned copper. This 
cleaning is best accomplished by dipping the copper sheet 
into dilute nitric acid, and then rinsing it with distilled 
water. Note the deposit of silver and get its properties. 
This deposit is called the silver tree. Spread the silver out 
on a hard surface and rub it with a knife blade, or other 
piece of metal, to bring out its luster. If you get consider- 






FIRST YEAR CHEMISTRY 293 

able powdered silver, collect it, squeeze the water out of it, 
put it on the cover of a porcelain crucible and fuse it down 
to a silver button, by applying a small blast lamp flame 
upon it. Write the equation for the replacement, assuming 
that copper nitrate is trje other product. 

Experiment 144. Reduction of silver chloride. Have 

ready a test tube rack with test tubes; also some silver ni- 
trate solution, hydrochloric acid, and mossy zinc. 

Precipitate some silver chloride from silver nitrate solu- 
tion by means of hydrochloric acid or a solution of any 
insoluble chloride. Shake and decant, leaving the clotted 
silver chloride in the test tube. Add to the precipitate a 
little dilute hydrochloric acid of a strength of 1 to 5. Of 
course, no action takes place. Then drop in two or three 
small pieces of mossy zinc. This reacts with the hydro- 
chloric acid, evolving hydrogen. Hydrogen, just at the 
moment of evolution and before it has had a chance to take 
on the definite form of a gas, is more active than when in 
the gaseous state; it is then called nascent hydrogen. Na- 
scent hydrogen is a powerful reducing agent, and in this ex- 
periment we have nascent hydrogen formed in contact with 
silver chloride. Under these circumstances, the silver chlo- 
ride is easily reduced to metallic silver, the hydrogen unit- 
ing with the chlorine. Note the powdery metallic silver, 
and bring out its luster as in the preceding experiment. 
"Write the equation for the production of nascent hydrogen 
from zinc and hydrochloric acid. Also write the equation 
for the reduction of silver chloride by nascent hydrogen. 
In this equation it is allowable to write hydrogen as H in- 
stead of H 2 , because it has not yet taken on the form of a gas. 

Experiment 145. Silver oxide; its preparation and its 
properties. Have ready a test tube rack with test tubes and 



294 FIRST YEAR CHEMISTRY 

a Bunsen burner; also some silver nitrate solution, stick 
sodium hydrate, and distilled water. 

In Experiment 139 you tried to form silver oxide, prob- 
ably without success. The compound is known, however, 
and may be made from silver nitrate. In Experiment 138 
you learned that metallic hydrates are generally made by 
precipitation from a solution of the salt by means of 
a solution of sodium hydrate. Silver is one of the very 
few exceptions to this rule. Sodium hydrate precipitates 
from a silver nitrate solution, not silver hydrate, but 
silver oxide. Try this experiment and get the properties 
of silver oxide. Write the equation for the precipitation, 
assuming that sodium nitrate is formed and stays in solu- 
tion. It may help in writing the equation, if you assume 
that silver hydrate is formed momentarily, but that it de- 
composes immediately into silver oxide and water. 

What effect did heat have upon lead nitrate? Do you 
think heat might have a similar effect on silver nitrate? 
Try this by evaporating about 10 c.c. of silver nitrate solu- 
tion just to dryness in a test tube and when the crystals 
of silver nitrate have formed, continue heating a little longer 
to see if the silver nitrate decomposes. Compare the color 
of the residue after heating with the color of the precipitate 
formed in the first part of this experiment. 

Experiment 146. Bismuth and its properties. Have 

ready a test tube and a Bunsen burner; also some lumps of 
bismuth. 

Examine a small lump of bismuth and get its properties, 
comparing it with other metals you have studied. 

Note on bismuth. — Bismuth usually occurs free in nature, tho the 
oxide and sulphide occur as minerals. It may be prepared from the 
native metal by melting it and allowing the melted metal to drain 
away from the impurities; or the ores may be roasted and reduced 



FIRST YEAR CHEMISTRY 295 

with carbon. The specific gravity of bismuth is 9.9; it melts at 265°C, 
and boils at about 1300°C. Its particular use is as an ingredient of 
fusible alloys. 

Experiment 147. Oxidation of bismuth. Have ready a 
brick, and the blast lamp; also some small lumps of bismuth. 

Put a small lump of bismuth on the brick and play a 
small blast lamp flame upon it till the bismuth catches fire. 
What is the color of the flame ? Get as many properties as pos- 
sible of the oxide. Write the equation for oxidizing bismuth. 

Experiment 148. The nitrates of bismuth ; their prep- 
aration and properties. Have ready a test tube rack with 
test tubes, and a Bunsen burner; also some bismuth, concen- 
trated nitric acid, and distilled water. 

First prepare bismuth nitrate by dissolving a small lump 
of bismuth in concentrated nitric acid. What is the color 
of the bismuth nitrate solution? What is the formula for 
bismuth nitrate, bismuth being a trivalent element? 

To a little of your bismuth nitrate solution in a test tube 
add water in very small portions and with constant shak- 
ing, till the white precipitate that is formed no longer re- 
dissolves. Then add still more water. Do you get more 
precipitate? Now add concentrated nitric acid, a few drops 
at a time, and shake till the precipitate is entirely dissolved. 
Again, add water till the precipitate appears; then add con- 
centrated nitric acid again. Does the precipitate redissolve? 
The white precipitate formed is called basic bismuth nitrate. 
The precipitation of this compound is the regular test for 
bismuth. The alternate formation and dissolving of this 
basic bismuth nitrate is called hydrolysis. 

Note on normal salts and basic salts. — Several of the heavy metals 
form basic salts, but most of the salts we have made and studied this 
year have been normal salts. A normal salt is a salt formed by replac- 



296 FIRST YEAR CHEMISTRY 

ing all the hydrogen of an acid by the metal, or, in other words, it is a 
salt in which the metal is united only to the acid radical; FeS0 4 (iron 
sulphate), CaCl 2 (calcium chloride), and Bi(N0 3 ) 3 (bismuth nitrate) 
are examples of normal salts. A basic salt is a salt in which the metal 
is united both to the hydroxyl group and to an acid radical; it may be re- 
garded as a normal salt in which part of the acid radical has been 
replaced by the hydroxyl group; Bi(OH)(N0 3 ) 2 , and Bi(OH) 2 (N0 3 ), 
are examples of basic salts. 

Note on hydrolysis. — This phenomenon depends upon the formation 
of basic salts. Hydrolysis is the formation of a basic salt by the addi- 
tion of water to a solution of the normal salt. The hydrolysis in the above 
experiment may be represented by the following equations: 

Bi(N0 3 ) 3 + H 2 = Bi(OH)(N0 3 ) 2 + HN0 3 
Bi(N0 3 ) 3 + 2H 2 = Bi(OH) 2 (N0 3 ) + 2HN0 3 
Bi(N0 3 ) 3 + 3H 2 = Bi(OH) 3 + 3HN0 3 

Bismuth chloride and antimony chloride are other well known salts 
which hydrolize. 

Experiment 149. Fusible alloy; its preparation and its 
properties. Have ready a clean deflagrating spoon, a test 
tube rack with test tubes, Bunsen burner, small beaker, 
tripod, gauze, horn-pan balance, set of smaller weights, 
and brass forceps; also some bismuth, tin, and lead. 

Weigh out exactly 15 grams of bismuth, 8 grams of tin, 
and 8 grams of lead. Mix the three metals and put them 
in a clean deflagrating spoon. Heat over a Bunsen burner 
till the mixture melts into one homogenous liquid, and then 
let it cool. When it has solidified, separate the alloy from 
the spoon, scratching off any adhering impurities, and put 
it in a small beaker. Add 25 c.c. of water to the beaker 
and bring to a boil. Does the alloy melt under the boiling 
water ? When it has melted pour the contents of the beaker 
(both water and metal) into a narrow thin walled test tube 
and let cool. What happens to the test tube as the alloy 
cools? How would you explain this effect of the solidifying 
alloy upon the test tube? 



FIRST YEAR CHEMISTRY 297 

To find out if any of the separate metals from which the 
alloy was made melt below 100°C, heat small lumps of bis- 
muth, of antimony, and of lead separately in boiling water. 
Does any one of them melt ? 

Note on alloys. — Many metals when fused together form a solid, 
homogenous, metallic-appearing mass upon cooling. This mass is 
called an alloy, and the constituents can not be separated by mechan- 
ical means. In general, the melting point of an alloy is below the 
average of the melting points of the constituents, and often it is lower 
than that of the lowest of them. Alloys can not be considered chemic- 
al compounds; they may be regarded as solid solutions of one metal 
in another. 

Some of the more important alloys are as follows: 

Common solder contains tin and lead. Fusible alloy contains bis- 
muth, lead, and tin. but the proportions vary with the use to which it 
is to be put; it is used as safety valves in steam boilers, and as fuse wires 
in electrical work. Type metal is mostly lead; the other metals are 
tin and antimony; it is used in making type. Pewter contains tin and 
lead. Brass contains copper and zinc, and is used for purposes where 
copper is wanted, and for which copper is too soft. Bronze varies 
considerably, but contains copper, tin, and zinc; it is used in coins, 
medals, statues, etc. Aluminium bronze consists of copper and alu- 
minium; it is used as an imitation of gold and in building hulls of yachts. 
Gun metal contains copper and zinc, and is used in fire-arms. Bell 
metal contains copper and zinc; bells are made of it. German silver 
consists of copper, zinc, and nickel; in appearance it resembles silver; 
its principal use is in electrical resistance coils; spoons, forks, and coins 
are sometimes made of it. Brittannia metal is an alloy of tin and 
antimony used in table ware. White metal contains less tin and more 
antimony than brittannia; it is used in machinery and ornaments. 
Babbit's metal varies considerably, but usually contains tin, copper, 
and antimony; it is used in bearings for machinery. Speculum metal, 
composed of copper and tin, is used in optical instruments. Valve 
metal contains copper, tin, and zinc. Silver coin is silver hardened 
with copper. Nickel coin is copper hardened with nickel. 

The accompanying table shows the approximate parts of each metal 
usually taken in making the alloy; the exact proportions vary consid- 
erably according to the use of the alloy. 



298 



FIRST YEAR CHEMISTRY 



Composition of Common Alloys 


Alloy 


Lead 


Tin 


Zinc 


Cop per 


Other Metals 


Common solder 


50 


50 








Fusible alloy- 


25 


25 






50 bismuth 


Type metal 


80 


10+ 






10 + antimony 


Pewter 


20 


80 








Brass 






30 


70 




Bronze 




4+ 


1+ 


95+ 




Aluminium bronze 








90 


10 aluminium 


Gun metal 






10 


90 




Bell metal 






25 


75 




German silver 






20 


60 


20 nickel 


Brittannia metal 




90 






10 antimony 


White metal 




90— 






10 + antimony 


Babbit's metal . 




89 




3 


8 antimony 


Speculum metal 




33 




67 




Valve metal 




10 


2 


88 




Silver coin 








10 


90 silver 


Nickel coin 


1 






75 


25 nickel 



Experiment 150. An experiment to prove the composition 
of ammonia. (Lecture Experiment.) Have ready a 4 liter 
prescription bottle and a two-hole rubber stopper to fit it, 
several pieces of glass tubing and of rubber tubing, two 
tubing clamps, pint fruit jar, triangular file, pneumatic 
trough, the generator for making hydrogen, long delivery 
tube, catch bottle, hard glass tube 15 cm. long and of about 
7 mm. bore, tripod, gauze, Bunsen burner and test tube; 
also some mossy zinc, copper turnings, sulphuric acid, nitric 
acid, platinum sponge and test papers. 

Thru one hole of the rubber stopper pass a straight piece 
of glass tubing about 15 cm. long, letting it pass well into 
the bottle. Thru the other hole let a piece of glass tube, 
bent at right angles, pass just thru the stopper. To the 



FIRST YEAR CHEMISTRY 299 

outer end of each tube attach a few centimeters of rubber 
tubing, each carrying a tubing clamp. 

Remove the fittings from the bottle, pour in two fruit 
jarfuls of water and mark with a file or a label the hight 
which the water reaches. Pour in five jars more of water 
and mark the new hight. Fill the bottle with water and in- 
vert it over the pneumatic trough. 

Generate the colorless oxide of nitrogen from copper and 
nitric acid as in Experiment 89. When all the air has been 
expelled from the generator see that the gas in the genera- 
tor is colorless,' allow the colorless oxide of nitrogen to pass 
into the large bottle till the water has fallen to the first 
mark, and then remove the generator. You now have in 
the large bottle two jarfuls of nitric oxide. Make hydrogen 
from zinc and sulphuric acid as in Experiment 32, and when 
the safety tube test shows that the hydrogen is free from 
air, pass hydrogen into the bottle containing the nitric ox- 
ide till the water falls to the second mark and then remove 
the generator. Y'ou have now added five jarfuls of hydro- 
gen. The mixture of two volumes of oxide of nitrogen and 
five volumes of hydrogen in the bottle should be colorless. 

Remove the bottle from the pneumatic trough, keeping 
the hand over the mouth of the bottle to prevent the re- 
maining water from escaping. Set the bottle upright, and 
insert the stopper and fittings with tubing clamps closed. 
Connect the straight glass tube by means of its rubber tub- 
ing to the water tap as shown in Fig. 81 Connect the 
bent glass tube to a catch bottle containing water, and to 
this connect a straight hard glass tube, 15 cm. long and of 
about 7 mm. bore, supported on a tripod and gauze, and 
having some platinum sponge packed loosely in the middle 
of it so as to fill the whole bore of the tube. Turn on the 
water and force the gases slowly thru the water of the catch 
bottle and note the bubbles in order to tell how fast the 



300 



FIRST YEAR CHEMISTRY 



gases are passing. Gradually force the mixed gases slowly 
over the platinum. When the gases have driven all the air 
from the catch bottle, and not before, heat the platinum. 
If the heat is applied before the air has all been washed out 
of the apparatus an explosion might occur in the catch 




Fig 81. Apparatus for proving the composition of ammonia. 

bottle. From what gases might this arise? Note that water 
is formed and deposited in the hard glass tube beyond 
the heated portion.' Where did the oxygen and hydrogen 
come from that formed this water? Note the formation of 
a new gag whose action on moist test paper can be seen by 
holding a piece of moist test paper in the gas that issues 
from the end of the hard glass tube. Let us call the new 
gas ammonia. Note its odor; also pass some of the new 
gas into a test tube containing a little water and note that 
the water dissolves the gas. Note that the liquid in the 
test tube has the same odor as the ammonia gas. Try the 



FIRST YEAR CHEMISTRY 301 

action on test papers of this water solution, which is called 
aqua ammonia ; also try the action of the aqua ammonia 
from the shelf on test papers. 

What were the two gaseous factors started with? All 
the oxygen of the oxide of nitrogen formed a new oxide. 
What oxide? Was the remaining gas pure nitrogen? If 
not. what simple substances must there have been in the 
new gas? The valence of nitrogen in ammonia is 3. What, 
then, is the formula for ammonia? Write the equation to 
show the change. Tell from the equation how many vol- 
umes of steam and how many volumes of ammonia gas were 
produced from the two volumes of nitric oxide and five 
volumes of hydrogen. 

Note on the composition of ammonia. — The composition of ammonia 
gas has been proved by the above experiment. It may also be proved 
by either of the two following methods: 

(1) Pass dry ammonia gas (made by warming strong aqua am- 
monia) over red hot copper oxide in a hard glass tube. The equation 
for this reaction is: 

2XH3 + 3CuO = 3Cu + 3H 2 + N 2 . 

The water is found in the tube, and the gas evolved is recognized as 
nitrogen. Draw a sketch of the apparatus that you would use if you 
were required to do this experiment. 

(2) Pass dry nitrogen gas (made by burning phosphorus in air to 
remove the oxygen) over red hot powdered magnesium in a hard glass 
tube. A dirty-yellow substance, magnesium nitride, Mg 3 N 2 , results. 
The equation for this reaction is: 

N 2 + 3Mg = Mg 3 N 2 . 

Add water to the magnesium nitride and recognize the ammonia that 
i- evolved by its odor. The equation for this reaction is: 

Mg 3 X 2 + 3H 2 = 3MgO + 2NH 3 . 

You have already done this experiment in a rough way; perhaps you 
remember the odor of ammonia obtained by putting water on the 
magnesium oxide made from powdered magnesium. This is due to 



302 FIRST YEAR CHEMISTRY 

the fact that the magnesium powder takes not only oxygen from the 
air, but also nitrogen, forming a little magnesium nitride along with 
the magnesium oxide; magnesium is practically the only metal that 
does this. 



Experiment 151. Preparation of ammonium salts. Have 

ready three beakers, three evaporating dishes, tripod, gauze, 
Bunsen burner, and glass stirring rod; also some test papers, 
aqua ammonia, distilled water, and the three common acids, 
hydrochloric, sulphuric, and nitric. 

Ammonium and ammonia. — Ammonium, NH 4 , is a uni- 
valent radical that acts much like sodium, Na, and potas- 
sium, K, in that it forms an alkaline hydrate and a series 
of salts that, like the corresponding salts of sodium and of 
potassium, are white and soluble in water. The following 
list shows the correspondence between the ammonium salts 
and the sodium and potassium salts: 

NH4CI (Ammonium chloride) NaCl KC1 
(NH 4 ) 2 S0 4 (Ammonium sulphate) Na 2 S0 4 K 2 S0 4 
NH4NO3 (Ammonium nitrate) NaN0 3 KN0 3 

Ammonium has never been isolated as have been the metals, 
sodium and potassium, but it is said to have a metallic 
character. It is closely associated with ammonia, NH 3 , 
which is a colorless gas. It is sufficient for us to note that 
NH 3 dissolves in water, forming the alkaline ammonium hy- 
drate, NH 4 OH. The equation is: 

NH 3 + H 2 = NH 4 OH. 

Ammonium hydrate acts just like caustic soda and caustic 
potash, particularly towards acids. 

Take equal portions of about 10 c.c. each of aqua ammo- 
nia in three evaporating dishes. Dilute each with about 20 
c.c. of distilled water, and then neutralize each with an acid, — 



FIRST YEAR CHEMISTRY 303 

the first with hydrochloric acid, the second with sulphuric 
acid, and the third with nitric acid. Evaporate each to 
crystallization, and get the properties of ammonium chlo- 
ride, of ammonium sulphate, and of ammonium nitrate. 
Compare the salts you made with the ammonium salts on 
the shelf. Write the equation for each neutralization. 

Experiment 152. Preparation of ammonia from an am- 
monium salt. Have ready a test tube, small beaker, Bun- 
sen burner, and wash-bottle; also some ammonium chlo- 
ride, calcium hydrate, and turmeric paper. 

Mix a little dry crystallized ammonium chloride with an 
equal volume of dry powdered calcium hydrate in a test 
tube. Wipe off any calcium hydrate sticking to the glass 
at the mouth of the test tube. Add enough water to the 
mixture in the test tube to form a paste. Wet a piece of 
turmeric paper and stick it to the under side of a small 
beaker. Set the beaker over the mouth of the test tube 
and note the change in the turmeric paper. If the change 
does not take place readily, heat the test tube gently, then 
smell of the gas evolved in the test tube. What is the gas? 
Write the equation for the change, assuming that the other 
products are calcium chloride and water. This is the regu- 
lar test for ammonium. All ammonium salts act this way 
when treated with calcium hydrate and water. This is a 
good method for preparing ammonia on a small scale; it 
should be noted, however, that this method does not prove 
the composition of ammonia. 

Experiment 153. An experiment to illustrate the great 
solubility of ammonia in water. (Lecture Experiment.) 
Have ready a 250 c.c. flask, a long-necked Kjeldahl flask, 
corks to fit both flasks, several pieces of glass tubing and 
rubber tubing, ring stand and two clamps, test tube, Bun- 



304 



FIRST YEAR CHEMISTRY 



sen burner, a large generator, tripod, and gauze; also some 
aqua ammonia, hydrochloric acid, and a few lumps of litmus. 

Arrange the apparatus 
as shown in Fig. 82. The 
straight glass tube in the 
lower flask should reach 
nearly to the bottom of 
the flask. The tube in 
the upper flask should 
be drawn down to a tip 
of between 1 to 2 mm. 
bore, and the end of the 
tip should reach not 
quite the middle of the 
upper flask. 

Heat a few lumps of 
litmus in a test tube 
with water. Empty the 
blue solution into the 
250 c.c. flask and fill the 
flask about three quar- 
ters full of water, then 
add a few drops of acid 
to the blue solution till, 
after shaking, it just turns 
red. 
Remove the Kjeldahl flask from its support and fill it 
with ammonia gas made by heating 50 c.c. of concentrated 
aqua ammonia in the generator, and conducting the gas 
into the Kjeldahl flask by means of a rubber tube. When 
the Kjeldahl flask is full of ammonia gas, insert the cork 
and put the flask in connection with the lower flask. Blow 
a little in the right angle tube to start the action. The phe- 
nomenon seen in this experiment is called the ammonia 




Fig 82. The ammonia fountain. 



FIRST YEAR CHEMISTRY 305 

fountain. Why does the water change from red to blue 
when it gets into the Kjeldahl flask? What causes the 
water to run from the lower flask into the upper flask? 

Experiment 154. Nitrous oxide ; its preparation and its 
properties. Have read}' a 6 by 1 test tube, fitted with a 
cork, a short delivery tube, several smaller test tubes, a 
large beaker, tripod, gauze, Bunsen burner, and a ring stand 
with clamp; also some dry ciystallized ammonium nitrate. 

Nearly fill the large beaker with water, bring it to a boil, 
and then remove the flame. Put about 10 grams cf dry 
crystallized ammonium nitrate in the 6 by 1 test tube, insert 
the stopper, and heat cautiously. The salt melts and then de- 
composes at a rather low temperature. Avoid heating it 
too much, for at high temperatures it decomposes with ex- 
plosive violence. Catch the gas evolved in test tubes over 
the hot water in the beaker, for the gas is soluble in cold 
water. Test the gas with a glowing splinter. The gas, tho 
it rekindles the splinter, is not oxygen, but nitrous oxide, 
N 2 0; the other product is water. Write the equation for 
the decomposition of ammonium nitrate. 

Note on the oxides of nitrogen.— This is the third oxide of nitrogen 
you have made this year. There are two oxides more, making five 
in all. The complete list is: 

Nitrous oxide N 2 

Nitric oxide , NO 

Nitrogen trioxide N2O3 

Nitrogen dioxide N0 2 

Nitrogen pentoxide N2O5 

The nitric oxide and nitrogen dioxide you are already familiar with. 
The nitrous oxide is often called laughing gas ; and it is the gas often 
used in dentistry for producing unconsciousness. The nitrogen tri- 
oxide and nitrogen pentoxide are unstable compounds and are of no 
particular importance. They are sometimes called nitrous anhydride 



306 FIRST YEAR CHEMISTRY 

and nitric anhydride respectively, because with water they form ni- 
trous acid and nitric acid. 

Experiment 155. Preparation of silicic acid and silicon 
dioxide. Have ready a test tube rack with test tubes, glass 
stirring rod, and Bunsen burner; also some sodium silicate 
solution, and concentrated hydrochloric acid. 

Add a few drops of concentrated hydrochloric acid to 
some sodium silicate solution in a test tube. Note the pre- 
cipitated silicic acid, and get its chief properties. 

Decant as much of the clear liquid as possible and heat 
the precipitated silicic acid to dryness in the test tube. 
When it is ihoroly dry, empty it out on a piece of paper and 
note its gritty nature. 

Note on silicic acid and silicon dioxide. — Elementary silicon does 
not occur free in nature, but many of its compounds called silicates 
are found as rocks. About one quarter of the earth's crust is said 
to be silicon. Silicic acid probably has the formula H 4 Si04, silicon 
being a tetravalent element. This easily loses a molecule of water, 
leaving H^SiC^. The gelatinous precipitate obtained above probably 
consists of a mixture of both these acids. When they are heated to 
dryness, water escapes, leaving silicon dioxide, SiC>2. Write the equa- 
tion for the formation of silicon dioxide from each acid. Silicon di- 
oxide is very widely distributed in nature, as sand, quartz, amethyst, 
opal, flint, rock crystal, and jasper, the various colors being due to 
impurities. Sodium silicate probably has the formula Na2SiC>3; write 
the equation for making H^SiC^ from sodium silicate by means of hy- 
drochloric acid, assuming that the other product is sodium chloride. 

Experiment 156. Calcium silicate; its preparation and its 
properties. Have ready a test tube rack with test tubes, and 
a Bunsen burner; also some sodium silicate solution and 
some calcium chloride. 

To a little dilute calcium chloride solution in a test tube 
add a little sodium silicate. Note the precipitation of cal- 
cium silicate and get its chief properties. Write the equa- 



FIRST YEAR CHEMISTRY 307 

tion for the change, assuming that the formula for sodium 
silicate is Xa 2 Si0 3 , and that for calcium silicate is CaSiOs, 
also that sodium chloride is the other product. 

Note on glass. — Glass is an amorphous, transparent, sometimes trans- 
lucent solid; a mixture of silicates, one metal always being an alkali. 
The metals generally found in glass are potassium, sodium, calcium, 
and lead. Glass is divided into two kinds, (1) lime or crown glass, 
and (2) lead or flint glass. Lime glass is a mixture of silicates of cal- 
cium and sodium, or of calcium and potassium; it is cheaper than lead 
glass, harder, has a higher melting point, and is used for ordinary pur- 
poses as plate glass, windows, and chemical apparatus. Lead glass 
is a mixture of the silicates of lead and sodium or of lead and potas- 
sium; it is heavy, more expensive than lime glass, softer, lustrous, 
and is used for making lenses and cut glass. Glass is made by melt- 
ing sand with the proper amounts of quick-lime, red lead, sal soda or 
potash, according to the kind cf glass wanted. Bohemian glass is a 
potash lime glass much used in chemical apparatus on account of its 
hardness and resistance to chemical action. Glass is colored by add- 
ing different substances, — iron or chromium to get green glass, copper 
or cobalt to get blue glass, manganese and iron to get orange glass, 
and calcium fluoride to get white glass. 

Experiment 157. Aluminium silicate ; its preparation and 
its properties. Have ready a test tube rack with test tubes, 
and a Bunsen burner; also some aluminium sulphate and 
some sodium silicate solution. 

To a little dilute aluminium sulphate solution in a test 
tube add a little sodium silicate solution. Note the precipi- 
tation of aluminium silicate and get its chief properties. 
Write the equation for the change, assuming that sodium sul- 
phate is the other product. 

Note on aluminium silicate. — A little fairly pure aluminium silicate 
occurs in nature as kaolinite ; much aluminium silicate occurs com- 
bined with silicates of other metals in such minerals as mica, feldspar 
beryl, garnet, tourmaline, and clay. Kaolinite and clay are used in 
making: pottery. 

Note on pottery. — Pottery may be divided into the three classes, (1) 



308 FIRST YEAR CHEMISTRY 

Porcelain (or china), (2) Stone-ware, and (3) Earthenware. Porcelain 
is made by using a mixture of kaolinite, sand, and pure feldspar. When 
the mass cools it becomes hard and white. It is often glazed, i.e., 
coated with an easily fused silicate and then baked. Stone-ware is 
similar to porcelain, but it is coarser and made from less pure materials. 
Of it are made jars, jugs, and the cheaper grades of crockery. Earthen- 
ware is made from still cheaper and coarser silicates and is glazed by 
throwing salt into the baking oven just before the baking is com- 
pleted. Of it are made tiles, flower pots, drain pipe, ordinary bricks, 
fire-clay bricks, and crucibles. 



Experiment 158. Preparation of metallic sulphides by 
precipitation. Have ready a test tube rack with test tubes, 
Bunsen burner, glass stirring rod, and the generator for 
making hydrogen sulphide; also some zinc sulphate, lead 
nitrate, stannous chloride solution, stannic chloride solu- 
tion, iron sulphide, and sulphuric acid. 

First prepare, in separate test tubes, dilute solutions of 
zinc sulphate and of lead nitrate; in two other test tubes 
take small portions of stannous chloride (SnC^) solution, 
and stannic chloride (SnCU) solution. 

Do the rest of the experiment in the hood. 

Generate a slow stream of hydrogen sulphide and pass the 
gas into the test tube containing the zinc sulphate in the 
same way you passed hydrogen sulphide into copper sul- 
phate solution in Experiment 43. Note the color and form 
of the precipitated zinc sulphide and the rapidity with which 
it forms. Rinse off the tube delivering the hydrogen sul- 
phide with a little water, and then pass hydrogen sulphide 
into the lead nitrate solution. Repeat, using the two tin 
solutions, washing off the delivery tube each time before 
using it in a new solution. In each case note the color and 
form of the precipitate and the rapidity with which it forms. 
Write the equations for the four precipitations, remember- 
ing that stannous tin is bivalent, and stannic tin is tetra- 



FIRST YEAR CHEMISTRY 309 

valent. These four precipitations may be considered the test 
for zinc, the test for lead, and* the tests for tin, respectively. 

Ferrous and ferric iron. — Iron exists in two states, — the 
ferrous, or slightly oxidized state, and the ferric, or highly 
oxidized state. FeCl 2 , Fe(N0 3 ) 2 , and FeS0 4 are ferrous chlo- 
ride, ferrous nitrate, and ferrous sulphate respectively, while 
the corresponding ferric salts are FeCl 3 , Fe(N0 3 ) 3 , and 
Fe 2 (S0 4 ) 3 . The most convenient ferrous salt to use is FeS0 4 , 
and Fe(N0 3 ) 3 is a convenient ferric salt. Ferrous salts 
usually have a tendency to oxidize to the corresponding 
ferric salts on being exposed to the air, and the ferric salts 
may be reduced to ferrous salts by proper reducing agents. 
The reduction of ferric sulphate to ferrous sulphate and the 
oxidation of ferrous sulphate to ferric sulphate are consid- 
ered in the following paragraphs. 

Experiment 159. Reduction of an iron solution from the 
ferric state to the ferrous state. Have ready a small beaker, 
tripod, gauze, Bunsen burner, and glass stirring rod; also 
some crystallized ferrous sulphate, sulphuric acid, iron 
filings, and distilled water. 

Select good, clear crystals of ferrous sulphate, FeS0 4 . 
7H 2 0. Wash off with a little water as much as possible of 
the coating of ferric sulphate that always forms when FeS0 4 
stands for any length of time. Then dissolve the crystals 
in distilled water, heating a little if necessary. If the solu- 
tion is clear and has a green color, only ferrous sulphate is 
present. More likely the solution will be somewhat cloudy, 
due to the presence of some ferric sulphate, which does not 
dissolve as readily as the ferrous sulphate. Boiling the solu- 
tion generally changes the ferrous salt to the ferric salt, the 
solution becoming brownish and cloudy. If the solution is 
cloudy, add a little dilute H 2 S0 4 , and heat gently till the 



310 FIRST YEAR CHEMISTRY 

solution becomes clear and green. If the solution does not 
clear up readily, add a few clean iron filings; the nascent 
hydrogen evolved reduces the ferric sulphate to ferrous 
sulphate. 

Fe 2 (S0 4 ) 3 + 2H = 2FeS0 4 + H 2 S0 4 . 

Save a little of the resulting solution for Experiment 160, 
and a little for Experiment 161. 

Experiment 160. Oxidation of an iron solution from the 
i e errous state to the ferric state. Have ready a test tube rack 
with test tubes, Bunsen burner, and glass stirring rod; also 
some concentrated nitric acid, and some freshly reduced 
ferrous sulphate solution from the preceding experiment. 

To a little of the freshly reduced ferrous sulphate solution 
made in the preceding experiment, add a few drops of con- 
centrated nitric acid and heat a little. Under these condi- 
tions the nitric acid is reduced to NO and H 2 0, liberating 
some free oxygen, which oxidizes the iron from the ferrous 
to the ferric state; nitric acid is, therefore, considered a 
good oxidizing agent. If the ferrous solution was concen- 
trated, the resulting ferric sulphate solution may be dark 
red. In that case dilute with a little water and compare 
the color with that of the ferrous sulphate solution started 
with. 

6FeS0 4 + 3H 2 S0 4 + 2HN0 3 = 3Fe 2 (S0 4 ) 3 + 2NO + 4H 2 0. 
Save a little of the resulting solution for Experiment 161. 

Experiment 161. The hydroxides of iron; their prepara- 
tion and their properties. Have ready a test tube rack with 
test tubes, glass stirring rod, and Bunsen burner; also some 
sodium hydrate, and the ferrous sulphate solution and the 
ferric solution from the two preceding experiments. 



FIRST YEAR CHEMISTRY 311 

First prepare a dilute solution of sodium hydrate by dis- 
solving a small stick of sodium hydrate in two thirds of a 
test tube of water. Add a little of this sodium hydrate so- 
lution to some clear, green ferrous sulphate solution in one 
test tube, and a little to some clear, ferric sulphate solution 
in another test tube. In each case note the color, form, and 
solubility in water of the precipitate. In one case the pre- 
cipitate is ferrous hydrate, Fe(OH) 2 , and in the other, ferric 
hydrate, Fe(OH) 3 . Which precipitate changes its color 
easily on standing? What is the change in the color? 
Which precipitate would you consider the more stable? 
Write the equation for each precipitation. 

Separate a little of the ferric hydrate from the liquid, 
evaporate it to dryness in a test tube, and bake it in the 
Bunsen burner flame for several minutes. The resulting 
powder is iron rust, water being evolved. Write the equa- 
tion for the decomposition of ferric hydrate by means of heat. 

Note on the oxides and hydroxides of iron. — There are three oxides 
of iron. (1) Ferrous oxide, FeO, a black, unstable powder, (2) Ferric 
oxide, Fe2C>3, iron rust, sold as rouge for polishing purposes, and as 
Venetian red for red paint, and (3) Ferroso-ferric oxide, Fe304, the black 
oxide of iron made by heating iron. The oxides of iron are found as 
ores, and from these the iron is extracted. Ferrous hydrate has no 
uses; freshly precipitated ferric hydrate is used as an antidote in case 
of arsenic poisoning. 

Note on the kinds of iron. — In the note on iron under Experiment 14, 
it was stated that the three different kinds of iron, cast iron, wrought 
iron, and steel, depend upon the percentage of carbon present, and 
the actual percentages were given. 

Note on the extraction of iron. — The ores of iron were mentioned 
in the note on iron under Experiment 14. Metallic iron is obtained 
from its oxygen ores by roasting them in air and then heating them 
with coke in a blast furnace, which allows the melted iron to run off 
at the bottom in the form of pig iron or cast iron. In the extraction 
of iron from its ores it is customary to add a flux to the charge of iron 
ore and carbon. This flux unites with the impurities in the ore and 
forms a fusible glass called slag, or cinder, and thereby prevents the 



312 FIRST YEAR CHEMISTRY 

reduced iron from reuniting with the oxygen of the air which is being 
constantly blown thru the mass. Wrought iron is made from cast 
iron by burning out the impurities in a puddling furnace. Steel may 
be manufactured by the following methods: (1) The Crucible Process. 
In this process iron and carbon are packed in tight, fire-clay boxes, 
and heated for several days. The iron slowly absorbs carbon during 
the heating and becomes steel. (2) The Bessemer Process consists in 
burning out the impurities in cast iron by passing air thru the molten 
metal, and then adding just enough cast iron to give the desired pro- 
portion of carbon; the operation is carried on in a converter, a huge 
egg-shaped vessel, supported so that it can be rotated into different 
positions. (3) The Thomas-Gilchrist Process is a modification of the 
Bessemer process; the converter is lined with a mixture of lime and 
magnesia in order to take up the impurities of sulphur and phosphorus 
occurring in the cast iron started with. (4) The Siemans-Martin, or 
the Open Hearth Process. In this process cast iron and wrought iron 
in proper proportions are melted on a hearth with an oxidizing gas 
flame; when a test shows that the metal contains the desired propor- 
tion of carbon, ferro-manganese is added to the charge, and it is then 
poured into molds; this method yields a tough, elastic steel. 

Note on the extraction of metals from ores in general. — Most metals 
are found in nature as compounds called "ores" from which the metal 
is obtained by means of the proper reducing agent. Most ores are ox- 
ides, and a few are sulphides. Any metallic oxide mixed with carbon 
(coal), and heated, gives the metal and oxide of carbon. If the ore 
is a sulphide, it is first roasted or heated to a high temperature; the 
sulphur burns off as sulphur dioxide; the metal is changed to the ox- 
ide, and this may be reduced as before with carbon. The above meth- 
ods are used technically on a large scale. In the laboratory the car- 
bon may be replaced by hydrogen or an easily oxidized metal. 

Experiment 162. The reduction of copper oxide by means 
of hydrogen. Have ready a hard glass tube, 20 cm. long 
and of about 7 mm. bore, drawn out to a capillary and turned 
up as in Experiment 32, a generator for making hydrogen, 
catch bottles containing sulphuric acid, tripod, gauze, test 
tube, Bunsen burner, and several rubber connectors; also 
some powdered copper oxide, mossy zinc, and sulphuric acid. 

Set up the apparatus as shown in Fig. 83, spread some 



FIRST YEAR CHEMISTRY 



313 



copper oxide along the bottom of the hard glass tube in the 
part' nearest the catch bottles. Generate hydrogen from 
zinc and sulphuric acid, dry it by passing it thru concen- 
trated sulphuric acid in the catch bottles, and pass a slow 




Fig. 83. Apparatus for reducing copper oxide with hydrogen. 

stream over the copper oxide. Light the hydrogen at the 
capillary with the safety tube. 'When the hydrogen is burn- 
ing at the tip, heat the hard glass tube gently. Note what 
happens. What deposit forms in the hard glass tube be- 
tween the copper oxide and the tip? Whence came the 
oxygen to form this water? To what was the copper ox- 
ide reduced? Write the equation for the reduction. 

Note on the reduction by hydrogen in general. — The method just de- 
scribed may be applied to the reduction of many metallic oxides. It 
is a method particularly convenient for use in the laboratory. 

Experiment 163. Calcium phosphate; its preparation and 
its properties. Have ready a test tube; also some sodium 
phosphate, and some calcium chloride. 



314 FIRST YEAR CHEMISTRY 

To a dilute calcium chloride solution add a little sodium 
phosphate solution. Note the precipitation of calcium phos- 
phate and get its chief properties. 

Note on phosphoric acid and phosphates. — There are several acids 
containing phosphorus, but the one that interests us most in our pres- 
ent state of progress has the formula H3PO4. The normal salt, Na3P04, 
is known, but the acid salt, Na2HP04, is more common. The sodium 
phosphate just used was the acid salt. The commonest calcium phos- 
phate is CaHP04; it was the salt precipitated in the preceding experi- 
ment. Write the equation for precipitating CaHP04 from calcium 
chloride solution by means of Na2HPC>4, assuming that sodium chlo- 
ride is the other product. 

Note on acid salts. — An acid salt is one in which part of the hydrogen 
has been replaced by a metal. Phosphoric acid, sulphuric acid, and 
carbonic acid form acid salts. Some of them are of commercial im- 
portance. The most common one is acid sodium carbonate, sometimes 
called sodium bicarbonate, NaHC03; it is a fine, white powder, closely 
resembling dry soda. Its principal use is in cooking, and it is then 
called cooking soda, or baking soda. 

Experiment 164. Preparation of illuminating gas from 
soft coal. Have ready a hard glass ignition tube with cork to 
fit it, ring stand and clamp, Bunsen burner, and a short hard 
glass tube drawn out to a capillary tip; also some soft coal. 

Half fill the ignition tube with coarsely powdered soft 
coal and clamp it to the ring stand. Make sure that all 
joints are tight. Heat the ignition tube gently at first, and 
then gently increase the heat. When the air has all been 
driven out of the tube, apply a flame to the capillary tip. 
Does the gas catch fire? Examine the residue in the tube 
after cooling. 

Note on illuminating gas. — In the note on carbon under Experiment 
45, it was stated that when soft coal is heated in air-tight iron retorts, 
the gases are expelled, and gray-black, porous, shiny coke is left in 
the bottom, and gray, shiny, fine-grained, hard gas retort carbon is 
deposited on the top of the inside of the retort. It was also stated 



FIRST YEAR CHEMISTRY 315 

that when the gas is passed thru water, tar is deposited, and the cleaned 
gas is stored in gas holders for illuminating purposes. The gas thus 
made is called illuminating gas, or coal gas. The hydrogen in the 
coal passes off, partly as pure hydrogen and partly in combination 
with carbon as hydro-carbons, such as marsh gas, CH 4 , and olefiant 
gas, C9H4. The nitrogen in the coal passes off as ammonia. The 
ammonia, carbon dioxide, and sulphur compounds are regarded as 
impurities in the gas and are removed before the gas is sent to the 
consumer. The ammonia is removed by the water that removes the 
tar. The carbon dioxide and the sulphur compounds are removed 
by means of purifiers, large, shallow boxes containing iron oxide and 
lime. The resulting gas is ready for use for illuminating and heating 
purposes. 

Note on water gas. — Water gas is made by forcing steam thru a 
mass of red hot coal and mixing the gaseous product with hot gas ob- 
tained from oil. When the steam comes in contact with the hot car- 
bon, the carbon reduces the steam to hydrogen, forming carbon mon- 
oxide. This change may be expressed by the following equation: 
C + H2O = CO + H 2 . This mixture of hydrogen and carbon monoxide 
burns with a pale blue flame; so before it can be used for illuminating 
purposes it must be enriched with gases which are illuminants. This is 
accomplished by spraying the heated products with oil that contains 
a high percentage of carbon. When the water gas is burnt this carbon 
is rendered luminous and gives to the gas its illuminating power. 
Owing to the high percentage of carbon monoxide, water gas and gases 
containing it are considered poisonous. Illuminating gas, on the other 
hand, is not considered poisonous. 

Note on acetylene. — Of late years a new gas. acetylene, has been put 
on the market for illuminating purposes. This gas has the formula, 
C 2 H 2 . and is made from calcium carbide and water, according to 
the equation. CaC 2 + 2H 2 = C 2 H 2 + Ca(OH) 2 . It is a colorless gas 
with an offensive odor. When it burns it yields an intense white 
light. 

Note on Pintsch gas. — This gas is frequently used for lighting rail- 
way coaches, because, unlike ordinary illuminating gas, its illuminat- 
ing power is not decreased by the pressure to which the gas is sub- 
jected in the storage cylinders. It is made by spraying naphtha into 
highly heated retorts in order to decompose the naphtha into other 
hydro-carbons. The resulting gases are cleaned and stored in a man- 
ner similar to that used in the manufacture of ordinary illuminating 



316 FIRST YEAR CHEMISTRY 

Experiment 165. An experiment to determine the per 
cent of water of crystallization in crystallized barium chlo- 
ride. Have ready a tripod, pipe-stem triangle, clean porce- 
lain crucible, Bunsen burner, horn-pan balance, set of small- 
er weights, and brass forceps; also some dry crystallized 
barium chloride. 

Ge,t the exact weight on the horn-pan balance of a clean, 
dry porcelain crucible without the cover, weighing it to 
centigrams. Put in the crucible not less than 2.5 grams 
and not more than 3.0 grams of finely powdered, dry crys- 
tallized barium chloride and get the exact weight to centi- 
grams. See that the barium chloride has no chips or other 
foreign matter in it. Record thus: 

Weight of crucible and crystallized barium chloride = g. 

Weight of crucible alone = g. 



Weight of crystallized barium chloride = g. 

Put the crucible and contents on the pipe-stem triangle 
and heat with the Bunsen burner flame that just touches 
the bottom of the crucible for about 5 minutes or till there 
is no danger of the crystals snapping out, or of boiling over 
when fusing. Then turn up the flame and heat at full heat 
for 15 minutes. Cool slowly by heating with a low flame 
for a few minutes. Then take the flame away entirely. 
When the crucible is hand warm, weigh again accurately 
to centigrams. 

Weight of crucible and salt before heating = g. 

Weight of crucible and salt after heating = g. 



Weight of water of crystallization = g. 

Heat to constant weight by repeating the above heating 
as often as necessary in order to make sure that all the 



FIRST YEAR CHEMISTRY 317 

water of crystallization has been driven off. Calculate the 
percentage of water of crystallization in the salt thus: The 
weight of the water of crystallization is to the weight of the 
crystallized salt as x is to 100; x equals the per cent of 
water of crystallization in crystallized barium chloride. 

Note on barium and its compounds. — The element barium is not 
found free in nature but its compounds are abundant. The most im- 
portant minerals containing barium are heavy spar or barytes (barium 
sulphate), and witherite (barium carbonate). The element may be 
prepared by electrolyzing fused barium chloride. The specific gravity 
is 3.6: it melts at red heat; its boiling point has not been determined. 
Metallic barium has no use. Barium chloride is the most important 
soluble barium compound and is used extensively in the laboratory, 
and in making other barium salts. Barium nitrate is used in making 
"green fire." Barium sulphide is used in making luminous paint. 
Barium sulphate is used in making paper and as an adulterant in paints. 

Experiment 166. Tests for bases and acids in salts. Have 

ready a test tube rack with test tubes, Bunsen burner, 
platinum test wire, and generator for making hydrogen 
sulphide; also some hydrochloric acid, sulphuric acid, lime- 
water, barium chloride solution, and silver nitrate solution. 

Strictly speaking, a base is any soluble alkaline hydrate, 
such as sodium hydrate, potassium hydrate, and ammonium 
hydrate, but the meaning of the term has been extended so 
that metals and metallic oxides are also considered bases. 
For that reason we often speak of testing for a base, when 
we mean simply testing for the presence of a metal. When 
we speak of testing for an acid in a salt we mean simply 
showing the presence of some acid radical in the salt. 

In testing salts to discover their composition, we usually 
test first for the base and then for the acid. 

Test tube reactions, using solutions, are called wet tests. 
Tests by flame coloration are called dry tests. The best 
way to apply the flame coloraton test is to heat the plati- 



318 FIRST YEAR CHEMISTRY 

num test wire red hot, and then dip the hot wire into the 
salt; enough salt usually sticks to the wire to give a dis- 
tinct flame coloration when the wire is again held in the 
flame. 

Previous experiments have shown us that some of the 
common bases and acids may be proved present by the fol- 
lowing simple reactions: 

Sodium gives a yellow flame coloration. 

Potassium gives a violet flame coloration. 

Silver is precipitated from a silver solution as a white, 
curdy precipitate of silver chloride by means of hydrochloric 
acid. 

Arsenic is precipitated as yellow arsenic sulphide by means 
of hydrogen sulphide. 

Lead is precipitated as black lead sulphide by means of 
hydrogen sulphide. 

Calcium is precipitated as white calcium sulphate by 
means of sulphuric acid. 

The test for a carbonate is to add an acid to the salt and 
test the gas evolved for carbon dioxide by means of lime- 
water. 

The test for a sulphate is to add barium chloride solution 
and get a white precipitate of barium sulphate which is in- 
soluble in dilute hydrochloric acid. 

The test for a chloride is to add silver nitrate solution 
and get a white, curdy precipitate of silver chloride. 

Analysis of an unknown salt. — Get an unknown salt from 
the instructor and determine its composition by the above 
tests. 

To show the presence or absence of all the metals treated 
in this book and of all the common acids requires a com- 
plete scheme of Qualitative Analysis. Since we have not 
the time in a one year's course to touch upon this subject 
further than we have in this experiment, it will be much 



FIRST YEAR CHEMISTRY 319 

better to leave the study of Qualitative Analysis for another 
year and to use then a text that treats of that subject alone. 

Experiment 167. An experiment to illustrate Berthollet's 
Laws. Have ready a test tube rack with test tubes; also 
some silver nitrate solution, sodium carbonate, and dilute 
hydrochloric acid. 

It is often desirable to be able to predict whether a re- 
action between two factors will take place. A study of a 
large number of experiments led Berthollet to conclude that 
when the factors are in solution, if one of the possible prod- 
ucts is insoluble or volatile under the conditions of the ex- 
periment, the reaction in question will take place. Further 
investigation has shown that this law does hold true and 
the reaction runs till one of the factors is used up. The 
only cases that are not covered by this statement are the 
very few neutralizations of acids with soluble alkalis. Ber- 
thollet's laws may be stated as follows: When two factors can 
form a product that is insoluble or volatile under the conditions 
of the reaction, that insoluble or volatile product will be formed 
until one of the factors is used up. Let us try two simple ex- 
periments to illustrate this law. 

Fill a test tube about one quarter full of silver nitrate 
solution. Add 3 or 4 drops of dilute hydrochloric acid, 
and shake a little to clot the precipitated silver chloride. 
Then add a few drops more of dilute hydrochloric acid and 
shake. Continue adding hydrochloric acid in small por- 
tions as long as any precipitate forms. When no more pre- 
cipitate forms upon the addition of more hydrochloric acid, 
the silver has all been changed from soluble silver nitrate 
to insoluble silver chloride. Any test that we may apply 
to the clear supernatant solution will not reveal a bit of 
silver therein. 



320 



FIRST YEAR CHEMISTRY 



solution, add successive small portions of dilute hydrochlo- 
ric acid till no more gas is evolved. The carbonate has 
now all been destroyed, and no further tests that we might 
apply will reveal the presence of any carbon in the liquid. 

The principle involved in Berthollet's laws is of great as- 
sistance in separating bases from each other in such tests 
as those we studied in the preceding experiment. 



Experiment 168. An experiment to determine the solu- 
bility of a salt in water. (Four 
students should work together on 
this experiment.) Have ready 
4 large beakers, sand bath, 4 tri- 
pods, 4 Bunsen burners, 3 porce- 
lain evaporating dishes, 3 gauzes, 
a 10 c.c. pipette, mortar and 
pestle, glass stirring rod, gradu- 
ate, horn-pan balance, set of 
smaller weights, brass forceps, 
thermometer, and 30 cm. rule; 
also some ammonium chloride, 
and distilled water. 

Put just 100 c.c. of distilled 
water in the large beaker and heat 
it on the sand bath as shown in 
Fig. 84 till the water just begins 
to boil. In the meantime add 
slowly, and with constant stir- 
ring, finely powdered ammonium 
chloride as long as any will dis- 
solve. When the water is boiling, 

there should still be a little undissolved salt left in the 
bottom of the beaker. The solution is then said to be 
saturated with salt. Turn the flame low, and keep the 




Fig. 84. A sand bath in use. 



FIRST YEAR CHEMISTRY 



321 



solution at just 100°C. Get the exact weight of an evapo- 
rating dish. 

Turn out the flame under the sand bath. With the mouth 
over the upper end of the pipette, draw up into the pipette 
some of the ammonium chloride solution till the liquid stands 
a little abovG the 10 c.c. mark. Put the forefinger over 
the end of the pipette as shown in Fig. 85; by loosening 

the finger a little, let the liquid 
run out till the surface falls to the 
10 c.c. mark. Now run this 10 
c.c. of ammonium chloride solu- 
tion into the weighed evaporating 
dish; evaporate it on a gauze till 
it shows the first signs of spatter- 
ing, and then immediately put the 
Mk dish over the steam bath as de- 
^njjfcf scribed in Experiment 91 and 
^^/EZ£$& evaporate to dryness. Weigh the 
^^S-??S$ ^ s k anc ^ con tents and determine 
tfyff*. the amount of ammonium chlo- 
\< ride dissolved in the 10 c.c. of 
water at 100°C. 

During the evaporation of the 
10 c.c. of ammonium chloride 
solution, the solution on the sand bath has been cool- 
ing. When it has fallen to 90°C, again take out 10 c.c. 
with the pipette and evaporate it to dryness in a weighed 
evaporating dish, and determine the amount of ammonium 
chloride dissolved in 10 c.c. of water at 90°C. 

Continue cooling the ammonium chloride solution; take 
out 10 c.c. at 80°, 70°, and so on till the solution reaches 
ordinary temperature. Kvaporate each portion to dry- 
ness on the steam bath and determine the amount of am- 
monium chloride that will dissolve in 10 c.c. of water at 



Fig. 85. A pipette in use. 



322 



FIRST YEAR CHEMISTRY 



the different temperatures. From these amounts deter- 
mine the amounts of ammonium chloride that would dissolve 
in 100 c.c. of water at 100°C., 90°C, 80°C, etc. 

























y 























8 

















/ 






7 























b 























b 























4 























3 























2 









/ 














1 









/ 















10° 20° 30° 40° 50° 6 0° 70° 80° 90° 

Fig. 86. Plotting the curve of solubility of a salt. 

Plotting the curve of solubility for ammonium chloride. — 

Draw in the notebook a square 10 cm. on each edge. Di- 
vide each side into 10 equal parts and draw perpendicular 
and horizontal lines so that the large square shall be divided 



FIRST YEAR CHEMISTRY 



323 



into small squares of 1 cm. on a side. Let each horizontal 
division represent 10°C, and let each perpendicular division 
represent 10 grams of dissolved substance. Label these 







V^ 








1 










9 


, 










1 










b 























7 























b 









I 














b 







—i, 


r 














4 








u 


B 


Vl 
















6 





c/ 


















2 


y 


y 


















1 


C>X 





















10° 20° 30° 40° 50° 60° 70* 80° 90° 

Fig. 87. 

Curves of solubility of some common salts. A, potassium carbonate; 

B, sodium chloride; C, potassium nitrate; D, sodium sulphate. 

divisions as shown in Fig. 86. On the vertical line repre- 
senting 100°C, put a small cross or dot opposite the number 
of grams of ammonium chloride you found were dissolved 



324 FIRST YEAR CHEMISTRY 

in 100 c.c. of water at 100°C. On every 10° vertical line 
indicate in like manner the number of grams of salt dis- 
solved in 100 c.c. of water at that temperature. Finally 
draw a line thru all the points. The resulting line is called 
the curve of solubility for ammonium chloride. The curve 
shown in Fig. 86 is for an imaginary salt and not for 
ammonium chloride. 

In Fig. 86 are shown the curves of solubility of several 
salts. These have been determined in the same way we 
have just determined the curve of solubility of ammonium 
chloride. By means of these curves of solubility it is possi- 
ble to tell by inspection the number of grams of dissolved 
salt at any temperature between 0°C. and 100°C. 

Note on saturation and kindred topics. — In the paragraph on test- 
ing for properties under Experiment 5 the subject of solubility was 
presented in a most elementary manner. When considering the sub- 
ject more in detail it is customary to call the dissolved substance the 
olute ; the liquid in which the solute is dissolved is called the sol- 
sent; and the solvent containing the dissolved solute is called the 
volution. The concentration of the solution refers to the amount of 
solute dissolved in a given amount of solvent; a dilute solution con- 
sains a small amount of solute; a concentrated solution contains a 
targe amount of solute; a saturated solution contains as much solute 
as the given volume of solvent can dissolve at the given tem- 
perature. By the solubility of a substance is meant the ratio 
of the weight of the solute to the weight of the solvent 
lin a saturated solution of the substance. Roughly speaking, a 
solid is very soluble, soluble, slightly soluble, or insoluble. In more 
exact work the solubility is represented as the number of grams of 
the solid in 100 grams of the saturated solution at any given tempera- 
ture. Concentrating a solution is evaporating part of the solvent, 
generally by boiling. Evaporating to crystallization is concentrating 
a solution until it becomes saturated at the boiling temperature, after 
which crystals appear as the solution cools. Evaporating to dryness 
is removing all the solvent by evaporation. The degree of solubility 
of solids in water depends upon the substance itself and the tempera- 
ture of the water. No two solids have exactly the same degree of 



FIRST YEAR CHEMISTRY 325 

solubility. In most cases solubility increases with the rise of tempera- 
ture. Supersaturated solution: Occasionally crystals do not deposit 
readily from a solution that is being evaporated. For instance, if a 
hot saturated solution of sodium sulphate entirely free from undis- 
solved solute be allowed to cool slowly and at absolute rest, a clear 
solution may be obtained that at the lower temperature contains more 
of the salt in solution than is ordinarily present in the saturated 
solution of the salt at that lower temperature. If the solution be 
now shaken, or if a crystal of sodium sulphate be dropped into the 
solution, crystals form immediately and rapidly until only enough of 
the salt is left in solution to form a saturated solution at that lower 
temperature. 

The periodic system. — It is quite proper that we should 
close our work upon the Atomic Period with a considera- 
tion of the Periodic System of classification of the elements, — 
a system that has been of the greatest value in classifying 
the enormous mass of data obtained by experiments. 

In 1S64, Newlands, an English chemist, made an ar- 
rangement of the elements according to their atomic weights, 
and noted that certain natural families existed in this list. 

In 1869, Lother Meyer, a German chemist, published a 
classification of the elements far more extended and better 
arranged than Newland's classification. He noticed that 
when the elements were written in a list in the order of 
the increasing atomic weights, the list divided itself roughly 
into blocks of about eight elements each, the corresponding 
numbers of successive blocks showing similarity. 

In the same year, Mendeleeff, a Russian chemist, made 
practically the same observation and developed the table 
still further. He went so far as to formulate the Periodic 
Law ; The properties of any element are periodic functions 
of its atomic weight, i.e., the properties of the elements 
vary as their atomic weights change. 

The accompanying table represents Mendeleeff's classifica- 
tion of the elements as he revised it shortly before his death. 



326 



FIRST YEAR CHEMISTRY 







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FIRST YEAR CHEMISTRY 



327 



To understand the full significance of the table, one 
needs a much wider knowledge than is possible to gain in 
one year of the properties of the elements and their com 
pounds, together with the reactions that take place between 
the substances themselves. 

One use of the Periodic Law is in determining the proba- 
ble atomic weights and other properties of substances 
which are still unknown, but whose probable existence is 
indicated by gaps in the Periodic Table. It may be of in- 
terest to know that Mendeleeff himself predicted that an 



Properties of 


Properties of 


Ekasilicon as Predicted 


Germanium as Found 


by 


by 


Mendeleeff in 1871 


Winkler in 1886 


Atomic weight = 72.8 


Atomic weight = 73 


Specific gravity = 5.5 


Specific gravity = 5.49 


Probably metallic 


Ge is a metal 


Probably dirty white 


Grayish white in color 


Melts with difficulty 


Melting point = 900°C. 


Will form an oxide when heated 


Forms an oxide when heated in 


in air 


air 


Will have 2 atoms of oxygen in 


Formula of the oxide is GeC>2 


the oxide 




Specific gravity of the oxide =4.7 


Specific gravity of Ge0 2 = 4.702 


May be obtained from its ore by 


Ge is obtained from its ore by 


reduction with carbon or sodi- 


reduction with carbon or hy- 


um 


drogen 


Will form a chloride with 4 atoms 


Forms a chloride, GeCl4 


of chlorine 




The chloride will boil near 100°C, 


GeCl 4 boils at 86°C. 


probably lower 




Specific gravity of the chloride 


Specific gravity of GeCl 4 = 1.887 


probably 1.9 




The chloride probably a liquid 


GeCl4 is a liquid 


Will probably form a sulphide 


Ge forms a sulphide, GeS2 


Scarcely acted on by acid- 


Ge is not acted on by acids 



328 FIRST YEAR CHEMISTRY 

element, whose atomic weight was about 72, might be dis- 
covered to fill that gap in the table which is now filled by 
the element, Germanium. The accompanying table shows 
the properties predicted by Mendeleeff in 1871 for his un- 
known element, which he called Ekasilicon; the table also 
includes the properties actually found by Winkler in 1886 
when he discovered Germanium. 

End of the atomic period. — This consideration of the Peri- 
odic Law and its value in chemical research closes our con- 
sideration of the Atomic Period. 

Organic chemistry. — In the early part of the Atomic Period 
it was believed that all animal and vegetable substances 
could be made only thru the influence of some " life princi- 
ple." All substances connected with living beings were, 
therefore, called organic substances. In 1828 Woehler suc- 
ceeded in preparing one of the substances occurring in the 
human body from inorganic mineral compounds. Since that 
time very many so-called organic substances have been pro- 
duced in the laboratory without the aid of any "life princi- 
ple." It has been discovered that all so-called organic sub- 
stances contain carbon. It is more convenient for us, there- 
fore, to consider Organic Chemistry not as the Chemistry of 
those compounds produced by some life principle, as it was 
formerly defined, but rather as the chemistry of the carbon 
compounds. Tho the number of organic substances is vastly 
larger than the number of inorganic substances, they are 
nearly all composed of only a few elements, seldom more 
than five or six. All organic compounds contain carbon; 
most organic compounds contain hydrogen; many organic 
compounds contain oxygen; some organic compounds con- 
tain nitrogen; and a very few contain sulphur and other ele- 
ments. In practice, organic substances are tested simply 
for carbon and for hydrogen, since it is usually not necessary 



FIRST YEAR CHEMISTRY 329 

to have definite information in regard to the other simple 
substances that may be present. Nearly all common or- 
ganic substances burn with a flame, hence, it is an easy 
matter to show the presence of carbon and of hydrogen. 
When the organic substance burns, carbon dioxide and 
water are formed and the presence of this can be shown by the 
regular tests that we have used during the year. 

Some important organic compounds. — From the long list 
of organic compounds we shall choose for further consid- 
eration -the following: ordinary alcohol, wood alcohol, ether, 
acetic acid, glucose, starch, petroleum, and soap. 

Ordinary alcohol, C 2 H 5 OH, often called ethyl alcohol, to 
distinguish it from other alcohols, is a colorless, volatile 
liquid, with a burning taste and a pleasant odor. It is 
made by fermenting molasses and distilling the resulting 
mixture, whereby the alcohol vapor condenses. It is lighter 
than water, but mixes with water in all proportions. It boils 
at about 78°C. Ordinary alcohol contains a considerable 
proportion of water. Absolute alcohol, i.e., strictly pure 
alcohol, may be obtained by removing the water from the 
ordinary alcohol by means of pure caustic lime. 

Wood alcohol, CH 3 OH, generally called methyl alcohol, is 
a colorless or slightly yellow liquid, with a sharp, peculiar 
odor. It boils at about 66°C, and mixes with water in all 
proportions. It is more poisonous than ordinary alcohol, 
and for that reason its use is limited to dissolving gums and 
shellacs in the manufacture of varnishes. It is made by the 
dry distillation of wood. Denatured alcohol is ethyl alcohol 
containing about 10 per cent methyl alcohol and a small 
amount of benzol, the last rendering it unfit for beverage 
purposes without destroying its fuel value. 

Ether, (C 2 H 5 ) 2 0, is a very light, easily volatile, colorless 
liquid with a characteristic odor. It boils at 35°C. ; it is 
extremely inflammable; it does not mix with water as 



330 FIRST YEAR CHEMISTRY 

does alcohol. Its chief use is as an anaesthetic. It is made 
by distilling a mixture of ethyl alcohol and sulphuric acid. 

Acetic acid, HC 2 H 3 27 is one of the commonest organic 
acids. It is made by the dry distillation of wood or by 
fermenting vinegar. It is a colorless liquid with a sharp, 
disagreeable odor and strong acid reaction. Its principal 
use is in making white lead, making artificial vinegar, and 
preparing acids. 

Glucose, C 6 H 12 6 , often called grape sugar, is found in 
many fruits, particularly in grapes. It is one of many 
sugars, cane sugar and beet sugar being other varieties. It 
may be made by boiling starch with dilute sulphuric acid. 
The principal use of glucose is as a substitute for cane sugar. 

Starch is a fine white powder, generally obtained from 
corn and potatoes. The formula is sometimes given as 
C6H 10 O 5 , but there is still considerable doubt as to the ac- 
curacy of this formula. The principal uses of starch are 
as a food, in laundries, and in making glucose. 

Petroleum is an oily liquid issuing from the earth. It has 
no definite chemical formula, because it is a mixture of va- 
rious hydrocarbons, i.e., compounds containing carbon and 
hydrogen. When subjected to distillation a large number of 
compounds are obtained. Among these compounds may be 
mentioned naphtha, benzine, kerosene, lubricating "oil, 
vaseline, and paraffine. These substances are generally col- 
lected in the order named, naphtha coming off at the lowest 
temperature, and paraffine remaining at the end of the dis- 
tillation. 

Soap is a mixture of sodium and potassium salts of organic 
acids, mainly stearic, palmitic, and oleic acids. It is made 
by boiling sodium hydroxide or potassium hydroxide with 
fats; the latter may consist of stearin, tallow, lard, house 
grease, cottonseed oil, palm oil, or cocoanut oil. The fol- 
lowing equation is typical of the changes that take place: 



FIRST YEAR CHEMISTRY 331 

C 3 H 5 (C 1S H 35 2 ) 3 + 3NaOH = 3XaC ls H 35 2 + C 3 H 5 (OH) 3 

glyceryl sodium sodium glyceryl 

stearate hydroxide stearate hydroxide 

(stearine) (caustic) (soap) (glycerine) 

The caustic alkali and the fat are boiled together, and salt 
is added to make the soap separate from the mixture and 
rise to the top. The glycerine is drawn off and the resulting 
soap is cooled, cut. and dried. Potassium hydroxide pro- 
duces soft soap, and sodium hydroxide produces hard 
soap. 

Experiment 169. An experiment to prove the presence of 
carbon and of hydrogen in certain organic compounds. Have 
ready a ring stand and medium ring, gauze, evaporating dish, 
four large beakers and Bunsen burner; also some lime-water, 
alcohol, ether, paper, wooden toothpicks, and a candle. 

Pour a few c.c. of alcohol into the evaporating dish and 
set fire to it. Xote the color of the flame. Hold over the 
flame a large cold, dry beaker. Note the dimming of the 
beaker on the inside. The presence of this moisture indi- 
cates the presence of hydrogen in the compound, the hy- 
drogen having burned and formed oxide of hydrogen. Re- 
move the beaker from the flame and immediately put the 
hand over the mouth of the beaker. Turn the beaker right 
side up, and without removing the hand entirely from the 
mouth of the beaker, pour in a little clear lime-water and 
shake it around in the beaker. Note the cloudiness ap- 
pearing in the lime-water. This indicates the presence of 
carbon in the compound, the carbon having burned and 
formed carbon dioxide, which reacted with the lime-water, 
precipitating calcium carbonate. 

Proceed with ether, paper, wood, and the candle in a simi- 
lar manner and prove the presence of both carbon and hy- 
drogen in each of these substances. 



332 FIRST YEAR CHEMISTRY 

The Modern Period.— In one of the opening paragraphs on 
the Theory of Chemistry it was stated that the Modern 
Period lasted from about 1900 to the present time. Several 
of the men associated with this period were mentioned un- 
der the Atomic Period, e.g., Faraday and Moissan. The men 
whose work is closely associated with modern conceptions 
of matter and of chemical change are Arrhenius, Ostwald, 
Dewar, Curie, Ramsay, Raoult, Pfeffer, Van't Hoff and Gibbs. 

Faraday's law. — Under the Atomic Period it was stated 
that Faraday showed that when the same electric current 
passes thru separate solutions of metallic salts the amounts of 
the different metals deposited are in the same ratio as the equiva- 
lent or combining weights of those metals. This may be il- 
lustrated by the following experiment: Arrange two beak- 
ers, — one containing a solution of silver nitrate and two 
strips of silver, the other containing a solution of copper 
sulphate and two strips of copper, — connected with each 
other and with a battery in such a way that the electric 
current may pass first thru one beaker and then thru the 
other. It is customary to speak of the current as entering 
the solution thru the anode, or positive electrode, and leav- 
ing by the cathode or negative electrode. By electrodes we 
mean the terminals which come from a battery or other 
source of electricity and which dip into a liquid. As the 
current passes thru the solution some of the silver leaves 
the solution and deposits on the cathode and a correspond- 
ing amount of silver dissolves off from the anode to replace 
that taken out of the solution; similarly for the copper. If 
the silver and copper cathodes are weighed before and after 
the current has been allowed to run for some time, it will 
be found that the amount of silver deposited on the silver 
cathode and the amount of copper deposited on the copper 
cathode are proportional to the equivalent weights of the 
two metals, i.e., to those weights of these two metals that 



FIRST YEAR CHEMISTRY 333 

would enter into an equation containing both. In other 
words, the increase in weight of silver is to the increase in 
weight of copper as twice the atomic weight of the uni- 
valent silver is to the atomic weight of the bivalent copper. 

Svante Arrhenius (1859 ), a Swedish physicist and 

chemist, formulated the theory of electrolytic dissociation 
as an explanation of solution. This theory is treated in the 
following paragraphs. 

Solution and dissociation. — Any liquid that will dissolve 
a substance is called a solvent. Water is the most impor- 
tant solvent, tho alcohol, ether, carbon bisulphide, and 
other liquids are often used. Soluble substances are divided 
into two classes, — electrolytes and non-electrolytes. An 
electrolyte is a substance whose water solution will allow an 
electric current to pass thru it; acids, bases, and salts are in 
this class. A non-electrolyte is a substance whose water 
solution will not allow an electric current to pass thru it; 
sugar and many organic substances are in this class. By 
electrolysis is meant the chemical changes caused by the 
passage of an electric current thru a dissolved or fused sub- 
stance. 

It is not known definitely just what " solution " is, but 
there is good reason for believing that, when an electrolyte 
is dissolved in water, a part of the substance is broken down 
or dissociated into simpler parts called ions, the breaking 
down process being called dissociation. Thus, HC1 breaks 
down into H ions and CI ions, HNO3 into H ions and NO3 
ions, H2SO4 into H ions and S0 4 ions, and AgN0 3 into Ag 
ions and N0 3 ions. These ions are not identical with atoms; 
they are particles (sometimes atoms) highly charged with 
electricity, hydrogen and metals positively, and acid radi- 
eals negatively. An ion with its charge of electricity is 
usually indicated by using the symbol of the element and 
writing above it the positive or negative sign, according to 



334 FIRST YEAR CHEMISTRY 

the kind of electricity carried by the ion. The dissociation 
of the substances just mentioned is illustrated by the fol- 
lowing equations: 

+ — 
HC1 = H + CI 
+ — 
HN0 3 = H + N0 3 

+ + — 
H 2 S0 4 = H + H + S0 4 

+ — 
AgN0 3 = Ag + N0 3 

The dissociation of a salt increases with the dilution of the 
solution and it is practically complete in extremely dilute 
solutions. In strong or concentrated solutions the salt is 
only partly dissociated and the clear liquid then contains 
(1) water, (2) some undecomposed or undissociated salt, 
(3) positive ions of the metal, and (4) negative ions of the 
acid radical. 

Let us now consider what happens when an electric cur- 
rent passes thru a dissolved salt, i.e., thru a dissociated 
electrolyte. If we dip the two electrodes from a battery 
into a beaker containing an aqueous solution of copper chlo- 
ride, the ions of copper will pass in the direction of the cur- 
rent from the anode to the cathode, and will separate from 
the solution as metallic copper on the cathode. At the 
same time the ions of chlorine will pass in the opposite di- 
rection and will separate from the solution at the anode as 
gaseous chlorine. The ion that moves in the same direction 
as the electric current is called a cation, and the ion that 
moves in the opposite direction is called an anion. As Ar- 
rhenius pointed out, only those substances conduct electricity 
which are at least partly dissociated and, therefore, the con- 
ductivity is due to the dissociated parts. 



FIRST YEAR CHEMISTRY 335 

Application of electrolysis. — The principles just considered 
are at the basis of such industries as electrotyping, electro- 
plating, refining copper, making metallic magnesium, sodium, 
and aluminium, making sodium hydrate, and in extract- 
ing and purifying gold, silver, and lead. 

Wilhelm Ostwald (1853 ), a German chemist, has con- 
tributed greatly to the modern theory of solution. 

James Dewar (1S42 ), an English chemist, has done 

much work on the liquifaction of gases. This has opened 
up a wide field in that it is now possible by the use of lique- 
fied gases to study chemical changes at extremely low tem- 
peratures. 

Pierre Curie (1859-1906), a French physicist and chemist, 
and Mme. Marie Curie, discovered in uranium ores a new ele- 
ment, radium, that emits so-called "radium rays." This 
led up to the disintegration hypothesis of Rutherford and 
Soddy. 

William Ramsay (1852 ), an English chemist, during 

his study of atmospheric nitrogen, discovered argon, helium, 
neon, krypton, and xenon, thereby adding a whole new family 
to the Periodic Table as originally arranged by Mendeleef. 
His discovery that radium is decomposed to helium, has 
opened anew the question of the derivation of the elements 
from one fundamental substance as set forth by Prout. 

Francois Marie Raoult (1830-1901), showed the relation 
between boiling points, freezing points, and molecular weights. 
It is generally true that when a substance dissolves in water 
the boiling point of that solution is raised above 100°C, 
and its freezing point is depressed below 0°C. Raoult's 
Laws are: (1) When weights of substances that are propor- 
tional to ihiir molecular weights are dissolved in equal volumes 
of a solvent the same rise in the boiling point is caused in 
each ea.« ; and (2) When wi ights of substances that are propor- 
tional to their molecular weights are dissolved in eqibal volumes 



336 FIRST YEAR CHEMISTRY 

of a solvent the same lowering of the freezing point is caused 
in each case. These laws hold in all cases except electro- 
lytes. Their principal use is in determining atomic weights. 
Thermochemistry. — This branch of chemistry carries on 
the work begun by Dulong and Petit. Practically every 
chemical change is accompanied by an absorption or pro- 
duction of heat. Heat of reaction is the number of calories 
of heat set free or absorbed when molecular quantities of 
substances enter into reaction. A calorie is the amount of 
heat necessary to raise the temperature of 1 c.c. of water 
from 0°C. to 1°C. An exothermic reaction is one in which 
heat is set free. An endothermic reaction is one in which 
heat is absorbed. The following equations are examples of 
exothermic and of endothermic reactions: 

H 2 + Cl 2 = 2HC1 + 22,000 calories. 
C + S 2 = CS 2 — 19,600 calories. 

Pfeffer showed that osmotic pressure corresponds to the 
pressure of gases. Dissolved substances in solution exert a 
pressure which is analagous to gaseous pressure and which 
is called osmotic pressure. Osmotic pressure may be il- 
lustrated as follows. Place in a vessel of water a vertical 
tube closed at its lower end by a semipermeable membrane 
and containing a solution of sugar. Let the levels of the 
liquids be at the same hight. Soon the column of liquid 
in the tube begins to rise, water entering from the outer 
vessel thru the membrane. 

Jacobus Hendricus Van't Hoff (1852 ), a Dutch chem- 
ist, has contributed to the modern theory of solution. He 
formulated the following statement concerning osmotic pres- 
sure: At the same osmotic pressure and temperature equal 
volumes of all solutions contain the same number of molecules 
and, in fact, that number which under the same pressure and 
at the same temperature exist in the same volume of a gas. 



FIRST YEAR CHEMISTRY 337 

Reversible reactions. — We have seen that 2H 2 + 2 = 2H 2 
by burning hydrogen in air or oxygen, but we have also 
seen that 2H 2 — 2H 2 -+- 2 b} T electrolysis. One is just 
the reverse of the other, but the conditions of the experiment 
were different. This running in both ways is indicated. thus: 

2H 2 + 02^=^2H 2 

and the reaction is called a reversible reaction. Other ex- 
amples of reversible reactions are: 

CaO + H 2 0,=^Ca0 2 H 2 
H 2 + S^=^H 2 S 

All reactions, except neutralization and those reactions that 
run according to Berthollet's Laws, may be considered re- 
versible reactions. 

Chemical equilibrium. — This may be illustrated as follows: 
Treat a little water in a test tube with an excess of table 
salt. Heat and shake well. The salt goes into solution 
and continues to do so until the solution is saturated, i.e., 
has dissolved all it can. If the solution is then cooled, some 
of the salt in solution will crystallize out, because salts are 
generally more soluble in hot water than in cold. During 
the heating particles of salt are passing from the crystal 
state to the solution state, and during the cooling from the 
solution state to the crystal state. It is thought that at the 
moment the salt is originally put into the water, particles of 
the salt immediately pass into the solution state, and while 
more is going the same way, some of that already in solu- 
tion begins to fly back into the crystal state, but at not so 
great a rate; the effect is that the solution increases in 
strength until the number of particles of salt going from the 
crystal state to solution state is exactly equal to the number 
of particles going in the opposite direction, i.e., action con- 
tinues until equilibrium is reached. 



33S FIRST YEAR CHEMISTRY 

In a similar manner all reversible reactions are supposed 
to run in both directions till equilibrium is reached. 

The law of mass action. — The conception of reversible 
reactions and of equilibrium figure in the law of mass ac- 
tion : When a reaction has reached the state of equilibrium 
the ratio of the product of the concentrations of the factors to 
the product of the concentrations of the products is always 
equal to a constant. Let A and B be two factors which re- 
act to form the products C and D. Let a, b, c, and d repre- 
g ent the concentrations of the four terms. 

Then A + B^=^C + D 

a b c d 

and, according to the law just stated, K being a constant, 

_^- = K 
cd 

Josiah Willard Gibbs (1839-1903), an American physicist 
and chemist, formulated what has been called the Phase Rule. 
As von Meyer in his History of Chemistry says: "The concep- 
tion of chemical equilibrium was given definite expression by 
Willard Gibbs in the phase rule, which has proved a valuable 
guide in numerous experimental researches of recent years. 
This theorem is mainly of value in representing clearly on a 
diagram the results of experiment, but it cannot be discussed 
here; the reader is referred to the literature on the subject." 

Conclusion. — With this rapid glance at the modern theories of chem- 
istry let us stop this first year's work, remembering that, as was stated 
before, theories are not facts but are only attempted explanations of the 
" Why " of such facts, as we learned during our study of the properties 
of chemical substances and the reactions of these substances with each 
other. By leaving the subject with this thought uppermost we should 
feel that the year's work has been a definite and successful step for- 
ward in our mental journey, and that we are in fit condition to tackle 
more advanced scientific problems, if such a course of study be our lot, 
or if it be not, simply to apply the common sense principles of scien- 
tific thinking to our every day problems in whatever field we find them. 



OUTLINE OF FIRST YEAR'S WORK 



This outline is designed to give a bird's-eye view of the 
first year's work in Elementary Chemistry as developed in 
this book. The order of subjects treated is consistent with 
such a view and does not follow the order in which they 
were studied in the laboratory. Nearly every formula 
given in this outline represents a substance made or stud- 
ied, and nearly every equation represents an experiment ac- 
tually performed, but not every substance or reaction stud- 
ied during the year is represented by a formula or by an 
equation. The equations given are typical and should serve 
as reminders of similar reactions; each one should bring to 
mind a complete experiment with apparatus and attendant 
phenomena. 

The numbers in brackets refer to the pages in the preced- 
ing text where a full treatment of the subject in question 
may be found. 

I. CHEMISTRY deals with the composition of sub- 
stances, with changes of substance, and with phe- 
nomena attendant upon such changes. [183] 

II. MATTER. 

(a) Elements: — Substances that cannot be divided into 

simpler substances. [35] 

(1) Metals,— e.g., Fe, Ag. 

(2) Nonmetals, — e.g., C, P. 

(b) Compounds: — Substances made by the union of two 

or more simpler substances, — e.g., CaO, Mn0 2 , 
H/'0 3 . HG, ZuSn,. (S 2 . [36] 

339 



340 FIRST YEAR CHEMISTRY 

(c) Pseudo compounds : — Substances similar to com- 
pounds but of varying composition. 

(1) Amalgam: — A semi-chemically united mixture 

of mercury and a metal, — e.g., NaHg, [Na x Hg y ], 
AuHg, ZnHg, NH 4 Hg. [148] 

(2) Alloy : — A semi-chemically united mixture of two 

or more metals, — e.g., solder, Sn and Pb; fusi- 
ble alloy, Sn, Pb, and Bi. [297] 

III. ALLOTROPY. 

(a) Definition : — The existence of an elementary sub- 

stance in different forms which have markedly dif- 
ferent physical and chemical properties, perhaps 
caused by a different number of atoms in the mole- 
cule. [73] 

(b) Examples :— C, P, S, O. 

IV*, CHANGES. 

(a) Physical : — A change that does not affect the compo- 

sition of the substance, — e.g., flame-tests, melting. 
[37] 

(b) Chemical : — A change that affects the composition of 

the substance. [38] 

(1) Analysis. Two kinds. 

(a) Proximate analysis: — The breaking down of 

a compound into simpler parts. 
CaC0 3 = CaO + C0 2 

(b) Ultimate analysis: — The breaking down of a 

compound into the elements of which it is 
composed. [45] 

2H 2 = 2H 2 + 2 

(2) Synthesis : — The building up of a compound from 

simpler parts. [109] 

Fe + S = FeS 
CaO + H 2 = Ca(OH) 2 



FIRST YEAR CHEMISTRY 341 

(3) Substitution : — The replacement of a simple sub- 

stance in a compound by another simple sub- 
stance capable of reacting with the compound. 
[99] 

H 2 + Mg = H 2 + MgO 
2AgX0 3 + Cu = Cu(N0 3 ) 2 + 2Ag 

(4) Metathesis : — The interchange of two simple sub- 

stances. [146] 
K2SO4 + Ba(N0 3 ) 2 = BaS0 4 + 2KN0 3 

V. OXIDATION. 

(a) Definition :— Union of oxygen with another substance 

forming an oxide or a higher oxide. [55] 

(b) Examples : — 

4Na L + 2 = 2Na 2 
2Ca IL + 2 = 2CaO 
4Sb in - + 30 2 = 2Sb 2 3 

c IV - + 2 = C0 2 
2CO + 2 = 2C0 2 
2Na 2 + 2 = 2Na 2 2 

VL REDUCTION. 

(a) Definition : — The taking away of oxygen (or any sim- 

ple substance) from an oxide (or any compound). 
[68] 

(b) Reducing agents. 

(1) Carbon:— 

2PbO + C 
2As 2 3 + 3C 

(2) Hydrogen:— 

CuO + H 2 
2AgCl + 2H 

(3) An easily oxidized metal : — 

CO2 + Zn = CO + ZnO 

(4) A metal (reduces an acid) : — 



2Pb 
4As 


+ C0 2 
+ 3C0 2 


Cu + H 2 
2Ag + 2HC1 






342 FIRST YEAR CHEMISTRY 



Fe + H 2 S0 4 


= FeS0 4 + H 2 


3Cu + 8HNO3 


= 3Cu(N0 3 ) 2 + 2 NO + 4H 2 


Zn + 2HC1 


= ZnCl 2 + H 2 


(5) Heat:— 




2HgO 


= 2Hg + 2 


H 2 C0 3 


= H 2 + C0 2 


(6) Electricity:— 




2H 2 = 


= 2H 2 + 2 


WATER. 





VII. 

(a) Of crystallization: — Water that is chemically united 

to a salt when it crystallizes out from a water solu- 
tion, — e.g. 

Na 2 CO 3 .10H 2 O, salsoda. 

CaS0 4 .2H 2 0, gypsum. 

FeS0 4 .7H 2 0, green vitriol. 

CuS0 4 .5H 2 0, blue stone. 

ZnS0 4 .7H 2 0, white vitriol. 

MgS0 4 .7H 2 0, Epsom salt. 

Na 2 SO 4 .10H 2 O, Glauber's salt. 

K 2 A1 2 (S0 4 ) 4 .24H 2 0, alum. 

BaCl 2 .2H 2 0, 

(b) Of solution: — Water in which a substance is dis- 

solved, generally expressed by the symbol Aq, — 
e.g., HC1 + Aq, NaCl + Aq. 

(c) Of dilution: — Water with which a liquid is diluted, 

generally expressed by the symbol Aq, e.g., H 2 S0 4 
+ Aq. 

(d) Hard water: — Water containing lime. [173] 

(1) Temporarily hard: — Water which has dissolved 
calcium carbonate by virtue of some carbon di- 
oxide already in solution; Aq + C0 2 + CaC0 3 ; 
may be softened by boiling off the C0 2 , where- 
by the CaC0 3 drops out. 



FIRST YEAR CHEMISTRY 343 

(2) Permanently hard: — Water which has dissolved 
calcium sulphate; Aq + CaS0 4 ; may be soft- 
ened by distilling off the water. 

VIII. ACIDS. 

(a) Definition : — A substance containing hydrogen which 

can be replaced by a metal or by a group of atoms 
(like ammonium, NH 4 ) that acts like a metal. 

(b) Methods of making. 

(1) By the addition of water to certain oxides : — 
P 4 + 50 2 = 2P 2 5 3H 2 + P 2 5 = 2H 3 P0 4 

C + 2 = C0 2 H 2 + C0 2 = H 2 C0 3 

S + 2 = S0 2 H 2 + S0 2 = H 2 S0 3 

2S0 2 + 2 = 2S0 3 H 2 + S0 3 = H 2 S0 4 

(2) By direct union : — 

H 2 + S = H 2 S 
H 2 + Cl 2 = 2HC1 

(3) By metathesis : — 

FeS + H 2 S0 4 = H 2 S + FeS0 4 

2KNO3 + H 2 S0 4 = 2HN0 3 + K 2 S0 4 

2NaCl + H 2 S0 4 = 2HC1 + Na 2 S0 4 

2NaBr + H 2 S0 4 = 2HBr + Na 2 S0 4 

2KI + H 2 S0 4 = 2HI + K 2 S0 4 

CaF 2 + H2SO4 = 2HF + CaS0 4 

(c) Nomenclature: — 

H 2 C03 carbonic acid 

H 3 P0 4 phosphoric acid 

HCIO3 chloric acid 

HNO3 nitric acid 

H 2 S0 4 sulphuric acid 

H 2 S0 3 sulphurous acid 

H2S hydrogen sulphide 

HC1 hydrochloric acid 

HBr hydrobromic acid 



344 FIRST YEAR CHEMISTRY 

HI hydriodic acid 

HF hydrofluoric acid 

(d) Kinds of acids. 

(1) Haloid acids : — Those composed of hydrogen and 

a halogen, — e.g., HF. 

(2) Oxygen acids: — Those containing oxygen, — e.g., 

H 3 P0 4 . 

(e) Tests for the acids themselves. 

(1) For H 2 C0 3 :— heat drives off C0 2 . 

(2) For HN0 3 :— Cu evolves NO. 

(3) For H 2 S0 4 :— CaCl 2 + Aq or BaCl 2 + Aq gives 

white precipitate insoluble in dilute HC1. 

(4) For H 2 S: — rotten egg odor. 

(5) For HC1: — AgN0 3 + Aq gives white precipi- 

tate that clots on shaking. 

(6) For HBr: — heat decomposes it, giving Br 2 vapors. 

(7) For HI: — heat decomposes it, giving I 2 vapors. 

(8) For HF:— etches glass. 

(f) Tests for the presence of the acid radicals in salts. 

(1) For H 2 C0 3 :— HC1 evolves C0 2 . 

(2) For HN0 3 :— H 2 S0 4 evolves HN0 3 or N0 2 . 

(3) For H 2 S0 4 :— CaCl 2 + Aq or BaCl 2 + Aq gives 

white precipitate insoluble in dilute HC1. 

(4) For H 2 S:— H 2 S0 4 evolves H 2 S. 

(5) For HC1: — AgN0 3 + Aq gives white precipitate 

that clots on shaking. 

(6) For HBr:— H 2 S0 4 evolves HBr or Br 2 (brown). 

(7) For HI:— H 2 S0 4 evolves HI or I 2 (violet fumes). 

(8) For HF: — H 2 S0 4 evolves HF which etches glass. 

IX. BASES. 

(a) Definition : — Metals, metallic oxides, and metallic hy- 

drates. [317] 

(b) Examples: — 



FIRST YEAR CHEMISTRY 345 

(NH 4 )'- 



Xa 1 - 


K 1 Ag 1 - Ca" Al m - 


Xa 2 


K 2 Ag 2 CaO A1 2 3 


NaOH 


KOH Ca(OH) 2 Al(OI 


(c) Tests. 




in By 


flame : — The dry salt on a pig 



platinum wire 
gives:— [318] 

Na, yellow; K, violet; Li, red; 

Ba, yellow green; Sr, red; Ca, orange. 
(2) By "wet way" (precipitation): — [317] 
Pb from Pb solution as PbS 
Ag from Ag solution as AgCl 
As from As solution as As 2 S 3 
Ca from Ca solution as CaS0 4 

X. NEUTRALIZATION. 

(a) Definition: — The metathesis of an alkali and an acid 

giving a neutral result. [138] 

(b) Examples : — 

XaOH + HC1 = XaCl 4- H 2 
2XaOH + H 2 S0 4 = Na 2 S0 4 + 2H 2 
Ca(OH) 2 + 2HC = CaCl 2 + 2H 2 

Ca(OH) 2 + H 2 S0 4 = CaS0 4 + 2H 2 
NH4OH + HNO3 = NH 4 N0 3 4- H 2 

2NH4OH + H 2 S0 4 = (NH 4 ) 2 S0 4 + 2H 2 

XL SALTS. 

(a) Definition : — The substance formed from the replace- 

ment of the hydrogen in an acid by a metal or a 
group of atoms acting like a metal. [101] 

(b) Methods of making. 

(1) From an acid and a metal, — e.g. 

2HC1 + Zn = ZnCl 2 + H 2 

(2) From an acid and a metallic oxide, — e.g., 

H.SO, + MgO = MgS0 4 + H 2 

(3) From an acid and a metallic hydrate, — e.g., 



346 FIRST YEAR CHEMISTRY 

HC1 + KOH = KC1 + H 2 

(4) From an acid and a salt of a weaker acid, — e.g., 

2HC1 + Na 2 C0 3 = 2NaCl + H 2 + C0 2 

(5) From two solutions that will give a precipitate 

when they are put together, — e.g., 
CaCl 2 + 2AgN0 3 = 2AgCl + Ca(N0 3 ) 2 

(6) By direct union: — e.g., 

Cu + S = CuS 

(c) Kinds. 

(1) Normal salts: — Those in which all the hydrogen 

of the acid has been replaced by the metal, 
—e.g., Na 2 C0 3 , K 2 S0 4 , Na 3 P0 4 , Ca 3 (P0 4 ) 2 . 
[295] 

(2) Acid salts : — Those in which part of the hydrogen 

of the acid has been replaced by the metal, — e.g., 
NaHC0 3 , KHS0 4 , Na 2 HP0 4 , CaHP0 4 . [314] 

(3) Basic salts: — Those in which the metal is united 

both to the hydrate group and to an acid radi- 
cal,— e.g., Bi(OH) 2 N0 3 . [296] 

(4) Double salts : — Two normal salts chemically united 

and crystallized out together, — e.g., Alum 
K 2 S0 4 .A1 2 (S0 4 ) 3 . 24H 2 0. [288] 

(d) Nomenclature : — 

H 2 C0 3 gives carbonates 
H 3 P0 4 gives phosphates 
HC10 3 gives chlorates 
HN0 3 gives nitrates 
H 2 S0 4 gives sulphates 
H 2 S0 3 gives sulphites 
H 2 S gives sulphides 
HC1 gives chlorides 
HBr gives bromides 
HI gives iodides 
HF gives fluorides 






FIRST YEAR CHEMISTRY 347 

XII. TESTS FOR GASES. 

C0 2 : — burning splinter, — goes out. Ca(OH) 2 + Aq 

gives white precipitate. 
X 2 : — burning splinter, — goes out. Ca(OH) 2 + Aq gives 

no precipitate. 
CO:— burning splinter, — gas burns with a blue flame. 
H 2 : — burning splinter, — gas pops and burns with a 

blue flame. 
2 : — glowing splinter, — splinter rekindles. 
S0 2 : — colorless; characteristic odor. 
XH 3 : — colorless; characteristic odor; alkaline reaction. 
HC1: — colorless; characteristic odor; acid reaction. 
H 2 S: — colorless; rotten egg odor. 
HF: — colorless; etches glass. 
NO: — colorless; air turns it brown. 
X0 2 : — brown; characteristic odor; acid reaction. 
Cl 2 : — greenish yellow; bleaches. 

XIII. PROVE THE COMPOSITION OF 

(a) Water. 

(1) By electrolyzing water and testing for hydrogen 

and oxygen. [56] 

(2) By burning hydrogen and condensing the moisture 

formed. [60, 89] 

(b) Sulphuric acid : — By synthesis. [78] 

(c) Hydrogen sulphide : — By synthesis. [110] 

(d) Hydrochloric acid : — By synthesis. [126] 

(e) Nitric acid and niter: — By metathesis, etc. [175] 

(f) Ammonia. 

CI) By means of nitric oxide and hydrogen. [298] 

(2) By means of dry NH 3 gas and red hot CuO. [301] 

(3) By means of dry X 2 . ho1 Mg, and water. [301] 

(g) Organic substances : — Burn the substance. The pres- 

ence of hydrogen is shown by the condensation of 



348 FIRST YEAR CHEMISTRY 

moisture on a cold beaker. The presence of car- 
bon is shown by shaking up the gaseous products 
of combustion with lime-water to test for carbon 
dioxide. [331] 
(h) Air: — By phosphorus, etc. [39] 

XIV. LAW OF BOYLE. (LAW OF MARIOTTE.) 

(a) Statement : — The volume of a gas varies inversely as 

the pressure to which it is subjected. P:P' = V':V 
or PV = P'V. 

(b) Experiment to verify : — Measure the volume of the 

air enclosed in the short arm of a Boyle Tube by 
mercury at the same hight in both arms. Add 
mercury till level in long arm above the level in 
the short arm equals the hight of the barometer. 
New volume of enclosed air is half the original 
volume. [220] 

(c) Application : — Reducing gas volumes. 

XV. LAW OF DALTON. (LAW OF CHARLES.) 

(a) Statement: — The volume of a gas varies directly as 

the temperature on the absolute scale, i.e., V:V' = 
273 + t : 273 + V; or the volume of a gas meas- 
ured at 0°C. increases or decreases by ^ ¥ of its 
volume for every degree rise or fall in tempera- 
ture. 

(b) Experiment to verify: — Heat in boiling water a 250 

c.c. flask, fitted with a tube and tubing clamp. 
Close tubing clamp. Cool flask in crushed ice. 
Open tubing clamp under water. Measure water 
that runs in. From the capacity of the flask and 
the contraction in air from 100° to 0° calculate 
the amount that 1 c.c. of air expands from 0°C 
to 1°C. [236] 

(c) Application : — Reducing gas volumes. 



FIRST YEAR CHEMISTRY 349 

XVI. LAW OF CONSERVATION OF MASS. (INDE- 
STRUCTIBILITY OF MATTER.) [LAVOISIER.] 

(a) Statement: — The sum of the weights of the prod- 

ucts of a chemical change is exactly equal to the 
sum of the weights of the factors. 

(b) Experiment to verify: — Make solutions of definite 

weights of barium nitrate and potassium sulphate. 
Mix the solutions. Filter. Dry and weigh the 
precipitated barium sulphate. Evaporate the fil- 
trate to dryness and weigh the potassium nitrate. 
Compare the sum of the weights of the factors 
with the sum of the weights of the products. [219] 

(c) Application : — In stoichiometry. 

XVII. LAW OF DEFINITE PROPORTIONS BY 

WEIGHT. [PROUST.] 

(a) Statement : — Every distinct chemical compound has 

a fixed and unalterable composition. 

(b) Experiment to verify: — Neutralize 25 c.c. of a solu- 

tion of salsoda with hydrochloric acid. Evapo- 
rate to dryness and weigh the salt. Treat 25 c.c. 
of a solution of salsoda with an excess of hydro- 
chloric acid. Evaporate to dryness and weigh 
the salt. Compare the two weights of salt, and 
note that the weights are equal. [221] 

(c) Application : — In stoichiometry and quantitative 

analysis. 

XVIII. LAW OF MULTIPLE PROPORTIONS BY 

WEIGHT. [D ALTON, PROUST.] 
(a) Statement: — When varying quantities of one sub- 
stance join a fixed amount of some other substance, 
the varying amounts of the first bear to each other 
a ratio expressed in simple numbers, as 1:2, 1:3, 
2:3, or the like. [222] 






350 FIRST YEAR CHEMISTRY 

(b) Examples: — CO and C0 2 ; ratio is 1:2. 

S0 2 and S0 3 ; ratio is 2:3. 
PbO and Pb 3 4 ; ratio is 3:4. 

(c) Application : — Of theoretical interest. 

XIX. LAW OF DEFINITE PROPORTIONS BY VOL- 

UME. [GAY-LUSSAC] 

(a) Statement: — In any chemical change the relative 

volumes of the gaseous factors and products bear 
to each other a simple numerical ratio. [226] 

(b) Experiment to verify: — Electrolyze water and note 

that for every volume of oxygen produced two 
volumes of hydrogen are given off. 

(c) Application : — In stoichiometry, and of theoretical in- 

terest. 

XX. BERTHOLLET'S LAWS. 

(a) Statement: — When two factors can form a product 

that is insoluble or volatile under the given condi- 
tions, that insoluble or volatile compound will be 
formed till one of the factors is used up. [319] 

(b) Experiment to verify: — To a solution of NaCl add 

some AgN0 3 + Aq in small portions and with 

shaking till no more precipitate is formed. 

NaCl + AgN0 3 = AgCl + NaN0 3 . 

Also add HC1 in small portions to some CaC0 3 till 

no more gas is evolved. 

2HC1 + CaC0 3 = CaCl 2 + H 2 + C0 2 . 

(c) Application : — In qualitative analysis. 

XXI. DALTON'S ATOMIC THEORY. [193] 

(a) Atoms: — Every simple substance is made up of mi- 
nute atoms, all alike, all of the same weight, and 
each one the smallest particle of that simple sub- 
stance that can enter into combination with other 
atoms, — e.g., S, Zn. 



FIRST YEAR CHEMISTRY 351 

(b) Molecules : — Every compound substance is made up 

of minute molecules, all alike, all of the same 
weight, and each a collection of atoms of diffeient 
simple substances grouped in simple and unalter- 
able numerical ratio and chemically united, — e.g., 
XaBr, PbCl 2 . Iv*C0 3 . 

(c) The molecule is the smallest particle of matter that 

can exist by itself. In a compound it is made up 
of atoms of different kinds, — e.g., NaCl. In a 
simple substance it is made up of atoms of the same 
kind, e.g., P 4 ,0 2 . Often the molecule of a simple 
substance consists of only one atom, — e.g., Fe, S. 

XXII. AYOGADRO'S SUGGESTION. 

(a) Statement : — Equal volumes of all substances, when 

in the state of gas, and under like conditions, con- 
tain the same number of molecules. [195] 

(b) Application : — In proving that the molecules of oxy- 

gen, of hydrogen, and of chlorine each have two 
atoms. [209] 

XXIII. DETERMINATION OF THE SPECIFIC GRAVI- 

TY OF GASES. 

(a) Definition of specific gravity : — The number of times 

greater that the weight of a definite volume of a 
substance is than the weight of an equal volume of 
some substance taken as a standard. Water is 
the standard for liquids and solids, and either air 
or hydrogen for gases. 

(b) Experiment to illustrate : — Get the capacity of flask 

and fittings; reduce this volume to S. T.. P. and 
find the weight of air in flask. AYeigh flask and 
fittings. Fill the flask and fittings full of the gas, 
and weigh again. From increase or decrease in 
weight and the weight of air the flask held get the 



352 FIRST YEAR CHEMISTRY 

weight of the gas. The specific gravity of the 
gas equals the weight of the gas divided by the 
weight of an equal volume of air under like con- 
ditions. [244] 
(c) Application : — In determining molecular weights. 

XXIV. DETERMINATION OF COMBINING NUMBER. 

(a) Definition of combining number: — The number ex- 

pressing the proportion by weight in which a sub- 
stance joins other substances. 

(b) Experiment to illustrate: — Treat a known weight of 

zinc with an excess of hydrochloric acid. Get 
the volume of hydrogen caught, reduce it to S. T. P. 
and get the weight of the hydrogen. The gram 
weight of the hydrogen caught is to the gram 
weight of the zinc used as 1 gram of hydrogen is 
to the gram weight of zinc necessary to produce 
1 gram of hydrogen. This last is the combining 
number of zinc. [247] 

(c) Application : — In determining atomic weights. 

XXV. DETERMINATION OF ATOMIC WEIGHTS. 

(a) Definition of atomic weight: — The number of times 

heavier an atom of an element is than an atom of 
hydrogen. 

(b) Method of determining : 

(1) Multiply the combining number by the valence. 

(2) From the molecular weight of a compound con- 

taining the element. 

(3) Divide 6.4 by the specific heat to find which mul- 

tiple of the combining number to take, because 
the specific heat multiplied by the atomic weight 
gives about 6.4. [Law of Dulong and Petit.] 

XXVI. DETERMINATION OF MOLECULAR WEIGHTS. 
(a) Definition of molecular weight: — The number of 



FIRST YEAR CHEMISTRY 353 

times heavier a molecule of a substance is than an 
atom of hydrogen, 
(b) Method of determining : — 

(1) By the physical method (used for gases) : — The 

molecular weight is twice the Sp. Gv. referred 
to hydrogen, because the Sp. Gv. is referred to 
the molecule of hydrogen, (H 2 ), while molecular 
weights are referred to the atom of hydrogen, 
(H). [256] 

(2) By the chemical method (used for liquids and 

solids) : — Change a chemical substance whose 
molecular weight is known into one whose mo- 
lecular weight is unknown, or vice versa. 
The molecular weight of the known is to the 
molecular weight of the unknown as the gram 
weight of the known is to the gram weight of 
the unknown. [257] 

XXVII. STOICHIOMETRY. 

(a) Definition : — Having given the gram weight of one 

substance that enters into a reaction, to determine 
the gram weights of the other members. [259] 

(b) Rule: — The molecular weight of the known is to the 

molecular weight of the unknown as the gram 
weight of the known is to the gram weight of the 
unknown. [261] 

(c) Application : — To determine the proper amounts of 

factors or products when working on a large scale. 

XXVIII. MODERN CONCEPTIONS. 

(a) Electrolysis : — Chemical changes caused by an elec- 

tric current in a dissolved or fused substance. 

(b) Electrolyte : — A substance whose aqueous solution 

conducts electricity. 



354 FIRST YEAR CHEMISTRY 

(c) Non-electrolyte: — A substance whose aqueous solu- 

tion does not conduct electricity. 

(d) Anode : — The positive electrode or pole by which the 

current enters the electrolyte. 

(e) Cathode: — The negative electrode or pole by which 

the current leaves the electrolyte. 

(f) Dissociation : — The breaking down into ions of a sub- 

stance when dissolved in water. 

(g) Anion : — The ion which moves against the current of 

electricity to the anode. 

(h) Cation: — The ion which moves with the current of 
electricity to the cathode. 

(i) Faraday's Law: — The quantities of the substances, 
separating at the electrodes in the same time, are 
in the proportion of their equivalent or combining 
weights. 

(j) Reversible reaction: — One that runs in one direction 
under certain conditions but in the opposite direc- 
tion under changed conditions. 

(k) Equilibrium : — All reactions are thought to be reversi- 
ble and to run in both directions till equilibrium 
is reached. 

(1) Heat of reaction : — Heat generated or absorbed dur- 
ing a chemical change, 
(m) Osmotic Pressure : — Pressure exerted by a substance 
in solution. 

(n) Law of mass action :— [338] 

(o) Phase rule:— [338] 



APPENDIX 



The Appendix contains a Table of the Elements with 
some of their Properties, Additional Problems for Class 
Use, Practical Questions, List of Text-books, Books for 
Reference Library, Lists of Apparatus and Chemicals 
needed for use with this book, and Suggestions to the 
Teacher. 

Table of the elements and some of their properties. — This 
table contains the complete list of the elements now known. 
The properties given and the authorities therefor are as 
follows: The exact atomic weights are taken from the 
Report of the International Committee on Atomic Weights 
for 1909, as printed in the January, 1909, number of the 
Journal of the Americal Chemical Society. The approxi- 
mate atomic weights are copied from the list on page 255 
of this book; they are in nearly every case the nearest 
whole number to the exact atomic weight. The valence, 
date of discovery and name of discoverer are taken from 
Buchka's Tables in Dammer's Handbuch der anorganischen 
Chemie. The state and color are taken from the author's 
Tables of Properties. The specific gravities, melting points, 
and boiling points are taken, as far as possible, from Bie- 
dermann's Chemiker-Kalendar for 1909. The other parts 
of the appendix explain themselves. 

355 



356 



FIRST YEAR CHEMISTRY 











The Elements and Some 


Name of 
element 


Sym 
bol 


Exact 
Atomic 
Weight 

0=16 


Ap- 
proxi- 
mate 
Atom- 
ic 
Weight 


Valence 


Specific Gravity 


Water = 1 


Air=l 


Aluminium 


Al 


27.1 


27 


III (?) 


2.6 




Antimony 


Sb 


120.2 


120 


Ill or IV 


6.7 




Argon 


A 


39.9 


40 


O 




1.38 


Arsenic 


As 


75.0 


75 


III or V 


5.7 




Barium 


Ba 


137.37 


137 


II 


3.6 




Bismuth 


Bl 


208.0 


208 


III or IV 


9.9 




Boron 


B 


11.0 


11 


III 


2.7 




Bromine 


Br 


79.92 


80 


I 


3.2 




Cadmium 


Cd 


112.40 


112 


II 


8.6 




Caesium 


Cs 


132.81 


133 


I 


1.8 




Calcium 


Ca 


40.09 


40 


II 


1.6 




Carbon 


C 


12.00 


12 


IV 


1.7-3.5 




Cerium 


Ce 


140.25 


140 


III or IV 


6.6 




Chlorine 


CI 


35.46 


35.5 


I 


1.3 


2.45 


Chromium 


Cr 


52.1 


52 


II to VI 


6.9 




Cobalt 


Co 


58 97 


59 


II or III 


8.9 




Columbium 


Cb 


93.5 


94 


V 


7.1 




Copper 


Cu 


63.57 


63.5 


II 


8.9 




Dysprosium 


Dy 


162.5 


162 


? 






Erbium 


Er 


167.4 


167 


Ill 






Europium 


Eu 


152.0 


152 


? 






Fluorine 


F 


19.0 


19 


I 




1.31 


Gadolinium 


Gd 


157.3 


157 


? 






Gallium 


Ga 


69.9 


70 


III? 


5.9 




Germanium 


Ge 


72.5 


73 


IV 


5.5 




Glucinum 


Gl 


9.1 


9 


II 


2.1 




Gold 


Au 


197.2 


197 


I or III 


19.3 





FIRST YEAR CHEMISTRY 



357 



of the 


ir Properties 










State 


Color 


Melting 
Point 
in t° C 


Boiling 
Point 
in t° C 


Date of 
Discovery 


Name of 
Discoverer 


Solid 


Gray 


C57 


1470-1700 


1827 


Woehler 


Solid 


Bluish-white 


430 


1500-1700 


1460 


Valentine 


Gas 


Colorless 


-189.5 


-185 


1894 


Rayleigh and Ram- 
say 


Solid 


Steel-gray to black 






13th cent. 


Albertus Magnus 


Solid 


Yellow 


850 


950 


1808 


Davy 


Solid 


Reddish-white 


269 


1435 


15th cent. 


Valentine 


Solid 


Brown 


infusible 


3500 


1808 


Gay-Lu ssac and 
Thenard 


Liquid 


Brownish-red 


—7.3 


63 


1826 


Balard 


Solid 


Bluish-white 


315 


860 


1841 


Stromeyer 


Solid 


Grayish-white 


27 


670 


1861 


Bunsen and Kirch- 
hott 


Solid 


Light-yellow 


760 




1808 


Davy 


Solid 


Black 


3500 (?) 




Ancient 


? 


Solid 


Gray 


>960 




1839 


Mosander 


Gas 


Greenish-yellow 


—102 


-33.15 


1774 


Scheele 


Solid 


Gray 


1515 




1797 


Vauquelln 


Solid 


Steel-gray 


1530 




1735 


Brand 


Solid 


Steel-gray 


1950 




1801 


Hatschett 


Solid 


Yellowish-red 


1065 


2100 


Ancient 


? 


Solid 


? 






? 


? 


Solid 


9 






1843 


Mosander 


Solid 


? 






? 


? 


Gas 


Yellowish-green 


-223 


-187 


1886 


Moissan 


Solid 


? 






? 


? 


Solid 


Gray 


30 




1875 


Boisbaudran 


Solid 


Grayish-white 


900 


>1350 


1886 


Winkler 


Solid 


Light-gray 


>960 




1828 


Woehler 


Solid 


Yellow 


1066 




Ancient 


? 



358 



FIRST YEAR CHEMISTRY 



The Elements and Some 


Name of 
element 


Sym- 
bol 


Exact 
Atomic 
Weight 


Ap- 
proxi- 
mate 
Atom- 


Valence 


Specific Gravity 










= 16 


ic 
Weight 




Water = l 


Air=l 


Helium 


He 


4.0 


4 







0.13 


Hydrogen 


H 


1.008 


l 


l 




0.07 


Indium 


In 


114.8 


115 


ill (?) 


74 




Iodine 


I 


126.92 


127 


I to Vll 


49 




Iridium 


Ir 


193.1 


193 


II to VIII 


21.2 




Iron 


Fe 


55.85 


56 


11, IV or VI 


7.0-7 8 




Krypton 


Kr 


81.8 


82 









Lanthanum 


La 


139.0 


139 


III 


6.1 




Lead 


Pb 


207.10 


207 


II or IV 


11.3 




Lithium 


Li 


7.00 


7 


I 


0.59 




Lutecium 


Lu 


r,4.o 


174 


9 


? 




Magnesium 


Mg 


24.32 


24 


II 


1.7 




Manganese 


Mn 


54.93 


55 


II to VIII 


7.2 




Mercury 


Hg 


200.0 


200 


II 


13.6 




Molybdenum 


Mo 


96.0 


96 


11 to VIII 


9.0 




Neodymium 


Nd 


144.3 


144 


III 






Neon 


Ne 


20.0 


20 









Nickel 


Ni 


58.68 


59 


II or IV 


8.7 




Nitrogen 


N 


14.01 


14 


III or V 




96 


Osmium 


Os 


190.9 


191 


11 to VIII 


22.5 




Oxygen 





16.00 


16 


II 




1.10 


Palladium 


Pd 


106.7 


107 


II or IV 


11.4 




Phosphorus 


P 


31.0 


31 


III or V 


yellow 1.8 




Platinum 


Pt 


195.0 


195 


II or IV 


21.5 




Potassium 


K 


39.10 


39 


I 


87 




Praseodymium 


Pr 


140.6 


141 


III 






Radium 


Ra 


226.4 


226 


? 






Rhodium 

i 


Rh 


102 9 


103 


II or IV 


12.1 





FIRST YEAR CHEMISTRY 



359 



of the 


ir properties 










State 


Color 


Melting 
Point 
in t° C 


Boiling 
Point 
in t° C 


Date of 
Discovery 


Name of 
Discoverer 


Gas 


Colorless 


.... • 


-267 


1S95 


Ramsay 


Gas 


Colorless 




-252.5 


1766 


Cavendish 


Solid 


White 


176 


red heat 


1863 


Reich and Richter 


Solid 


Bluish-black 


114 


184 


1812 


Courtois 


Solid 


White or gray 


2500 




1802 


Tennant 


Solid 


Gray-white 


1050-1500 




Ancient 


? 


Gas 


Colorless 


-169 


—151.7 


after 1895 


Ramsay 


Solid 


Grayish-white 


810 




1839 


Mosander 


Solid 


Bluish-white 


327 


1400-1600 


Ancient 


? 


Solid 


White 


186 


950 


1807 


Davy 


? 


o 


? 


? 


o 


? 


Solid 


Grayish-white 


750 


1100 


1830 


L i e b i g and 
Bussy 


Solid 


Gray-white 


1900 




1807 


Gahn and John 


Liquid 


Gray 


-39.4 


357 


Ancient 


? 


Solid 


Gray 






1790 


Hjelm 


Solid 


White 


840 




1885 


Welsbach 


Gas 


Colorless 






after 1895 


Ramsay 


Solid 


Grayish-white 


1484 




1751 


Cronstedt 


Gas 


Colorless 


—210.5 


—196 


1772 


Rutherford 


Solid 


Bluish-white 


2500 




1803 


Tennant 


Gas 


Colorless 


>-230 


—184 


(1774 
1 1775 


( Priestley 
1 Scheele 


Solid 


White or gray 


1900 




1803 


Wollaston 


Solid 


Yellow and red 


44.2 


290 


(1674 
1l676 


( Brand 
"1 Kunckel 


Solid 


Grayish-white 


1775 




1750 


Waston 


Solid 


Bluish-white 


62.5 


720 


1807 


Davy 


Solid 


White 


940 




1885 


Welsbach 


Solid 


? 






(1898) ? 


Curie 


Solid 


White 


2000 




1803 


Wollaston 



360 



FIRST YEAR CHEMISTRY 



The Elements and Some 


Name of 
element 


Sym- 
bol 


Exact 
Atomic 
Weight 


Ap- 
proxi- 
mate 
Atom- 


Valence 


Specific Gravity 










0= 16 


ic 
Weigh! 




Water = 1 


Air = 1 


Rubidium 


Rb 


85.45 


85 


I 


1.5 






Ruthenium 


Ru 


101.7 


102 


II or IV 


12.3 






Samarium 


Sa 


150.4 


150 


III 








Scandium 


Sc 


44.1 


44 


IV 








Selenium 


Se 


79.2 


79 


11 or IV 


4.5-6.5 






Silicon 


Si 


28.3 


28 


IV 


2.5 






Silver 


Ag 


107.88 


108 


I 


10.5 






Sodium 


Na 


23.00 


23 


I 


0.97 






Strontium 


Sr 


87.62 


88 


II 


2.5 






Sulphur 


S 


32.07 


32 


II or VI 


2.0 






Tantalum 


Ta 


181.0^ 


181 


V 


10.8 






Tellurium 


Te 


127.5 


128 


II or IV 


6.3 






Terbium 


Tb 


159.2 


159 


? 








Thallium 


Tl 


204.0 


204 


lor III 


11.9 






Thorium 


Th 


232.42 


232 


IV 


11 






Thulium 


Tm 


168.5 


168 


III 








Tin 


Sn 


119.0 


119 


II or IV 


7.3 






Titanium 


Ti 


48.1 


48 


II or IV 


3.5-4.9 






Tungsten 


W 


184.0 


184 


II to VI 


19.1 






Uranium 


U 


238.5 


239 


III or V (?) 


18.7 






Vanadium 


V 


51.2 


51 


Ill or V 


5.5 






Xenon 


Xe 


128.0 


128 











Ytterbium 


Yb 


172.0 


172 


III 








Yttrium 


Y 


89.0 


89 


III 








Zinc 


Zn 


65.7 


65.5 


II 


7.1 






Zirconium 


Zr 


90.6 


91 


IV 


4.2 







FIRST YEAR CHEMISTRY 



361 



of the 


ir Properties 










State 


Color 


Melting 
Point 
in to C 


Boiling 
Point 
in t° C 


Date of 
Discovery 


Name of 
Discoverer 


Solid 


White 


38.5 


696 


1861 


Bun sen and Kirch- 
boff 


Solid 


White 


>1950 




1845 


Claus 


Solid 


? 






1S79 


Koisbaudran 


Solid 


? 






1879 


Nilson and Cleve 


Solid 


Brown to black 


170-180 


690 


1817 


Berzelius 


Solid 


Brown 


1200 




1823 


Berzelius 


Solid 


White 


955 




Ancient 


? 


Solid 


Grayish- white 


95.6 


742 


1807 


Davy 


Solid 


Pale-yellow 


900 




1808 


Davy 


Solid 


Yellow 


115-119 


444.6 


Ancient 


? 


Solid 


Black (?) 


2250 




1802 


Eckeburg 


Solid 


Reddish- white 


446 


1390 


1782 


Reichenstein 


Solid ? 


? 






? 


? 


Solid 


White 


290 


1600-1800 


J 1861 
| 1862 


( Crookes 
1 Lamy 


Solid 


Greyish-white 






1828 


Berzelius 


Solid 


? 






1878 


Delafontaine 


Solid 


White 


232 


1450-1600 


Ancient 


? 


Solid 
Solid 
Solid 


Gray 

Steel-gray 

Gray 


3000 
1700 
800 




j 1791 
1 1795 
1 1781 
"I 1783 
\ 1789 
) 1840 


\ Gregor 

/ Klaproth 

I Scheele 

/J.&F.d'Elhujar 

( Klaproth 

1 Peligot 


Solid 


Gray 


1680 




1831 


Berzelius 


Solid 


Colorless 


—140 


—109.1 


after 1895 


Ramsay 


Solid 


? 






1878 


Marignac 


Solid 


Grayish-black 






1794 


Gadolin 


Solid 


Bluish-white 


419 


918 


16th cent. 


Paracelsus ? 


Solid 


Black 


1500 




1789 


Klaproth 



362 FIRST YEAR CHEMISTRY 

Additional Problems for Class Use 

The following additional problems are inserted here in 
case it is desirable for the class to have more practice than 
is afforded by the problems given in the body of the text. 
They are grouped under the headings: 

Metric System. 

Law of Multiple Proportions. 

Law of Boyle. 

Law of Dalton. 

Laws of Boyle and Dalton. 

Weight and Specific Gravity of Gases. 

Combining numbers. 

Molecular Weights by the Physical Method. 

Molecular Weights by the Chemical Method. 

Plain Stoichiometry. 

Stoichiometry involving Gas Volumes. 

Determination of Molecular Formulae. 

Determination of Percentage Composition. 

Problems on the Metric System 

1. Add 1 meter, 2 decimeters, 8 centimeters, and 5 millimeters 
and express the result in centimeters and decimal parts of a centi- 
meter. 

2. Add 10 grams, 8 decigrams, 5 centigrams, and 10 milligrams, 
and express the sum in grams and decimal parts of a gram. 

3. A certain length measures 46 centimeters. What is this length in 
millimeters? In decimeters? In meters? In inches? In feet? 
In reducing from the metric system to the English system use 2.5 cm. 
as the equivalent for one inch. 

4. A glass tube is 86 cm. long. How long is it in millimeters? 
In decimeters? In meters? In inches? In feet? In reducing from 
the metric system to the English system use 2.5 cm. as the equivalent 
for one inch. 

5. A box is 0.4 m. long, 2.5 dm. wide, and 12 cm. deep. What 
is its volume in cubic centimeters? 



FIRST YEAR CHEMISTRY 363 

6 A flask holds 625 c.c. What is its capacity in cubic millimeters? 
In cubic decimeters? In liters? In cubic meters? 

7. A flask holds 5S0 c.c. What is its capacity in cubic decimeters? 
In liters? 

8. The capacity of a bottle is 1540 c.c. How many liters does 
it hold? 

9. A room measures 6 m. long, 6 m. wide, and 3 m. high. If one 
c.c. of air weighs 0.00129 g., find the weight of air in this room, as- 
suming that the room is empty of everything except air. 

10. A recitation room is 12 m. long, 7.5 m. broad, and 5 m. high. 
Plow many cu. m. of space in the room when it is empty? How many 
liters in the room when it is empty? How many cu. m. of air does 
it contain when there are 30 persons present whose average volume 
is 75 cu. dm.? 

11. Alcohol is 0.8 as heavy as water. What is the weight of one 
liter of alcohol? 

12. A beaker holds 240 c.c. of water. How many grams of water 
does it hold? How many grams of glycerine can it hold, glycerine 
being once and a quarter as heavy as water? How many grams of 
sulphuric acid can it hold, sulphuric acid being 1.8 times heavier than 
water? 

Law of Multiple Proportions 

1. The two oxides of antimony have the following composition: 
Antimonious oxide contains 83.33% of antimony and 16.67% of oxy- 
gen; antimonic oxide contains 75.00% of antimony and 25.00% of 
oxysren. Calculate the simple ratio between the varying amounts 
of oxygen that unite with unit weight of antimony. 

2. The two sulphides of antimony have the following composition: 
Antimonious sulphide contains 71.42% of antimony and 28.58% of 
sulphur; antimonic sulphide contains 60.00% of antimony and 40.00% 
of sulphur. Calculate the simple ratio between the varying amounts 
of sulphur that unite with unit weight of antimony. 

3. The two oxides of arsenic have the following composition: 
The first oxide of arsenic contains 75.75% of arsenic and 24.25% of 
oxygen; the second oxide of arsenic contains 65.21% of arsenic and 
34.7')',' of oxygen. Calculate the simple ratio between the varying 
amounts of oxygen that unite with unit weight of arsenic. 

4. The two sulphides of arsenic have the following composition: 
I'll'- fir-t sulphide of arsenic contains 60.97% of arsenic and 39.03% 
of sulphur; the second sulphide of arsenic contains 48.38% of arsenic 



364 FIRST YEAR CHEMISTRY 

and 51.62% of sulphur. Calculate the simple ratio between the vary- 
ing amounts of sulphur that unite with unit weight of arsenic. 

5. The two oxides of barium have the following composition: 
Barium monoxide contains 89.54% of barium and 10.46% of oxygen; 
barium dioxide contains 81.06% of barium and 18.94% of oxygen. 
Calculate the simple ratio between the varying amounts of oxygen 
that unite with unit weight of barium. 

6. The two oxides of calcium have the following composition: 
Calcium monoxide contains 71.42% of calcium and 28.58% of oxygen; 
calcium dioxide contains 55.55% of calcium and 44.45% of oxygen. 
Calculate the simple ratio between the varying amounts of oxygen 
that unite with unit weight of calcium. 

7. The two chlorides of chromium have the following composition: 
Chromous chloride contains 42.27% of chromium and 57.73% of 
chlorine; chromic chloride contains 32.80% of chromium and 67.20% 
of chlorine. Calculate the simple ratio between the varying amounts 
of chlorine that unite with unit weight of chromium. 

8. The two oxides of hydrogen have the following composition: 
Water contains 11.11% of hydrogen and 88.89% of oxygen; hydro- 
gen peroxide contains 5.88% of hydrogen and 94.12% of oxygen. 
Calculate the simple ratio between the varying amounts of oxygen 
that unite with unit weight of hydrogen. 

9. The two chlorides of manganese have the following composi- 
tion: The first chloride of manganese contains 43.65% of manganese 
and 56.35% of chlorine; the second chloride of manganese contains 
27.91% of manganese and 72.09% of chlorine. Calculate the simple 
ratio between the varying amounts of chlorine that unite with unit 
weight of manganese. 

10. The two chlorides of mercury have the following composition: 
Mercurous chloride contains 84.92% of mercury and 15.08% of chlorine; 
mercuric chloride contains 73.80% of mercury and 26.20% of chlorine. 
Calculate the simple ratio between the varying amounts of chlorine 
that unite with unit weight of mercury. 

11. The two oxides of sulphur have the following composition: 
The first oxide of sulphur contains 50% of sulphur and 50% of oxy- 
gen; the second oxide of sulphur contains 40% of sulphur and 60% 
of oxygen. Calculate the simple ratio between the varying amounts 
of oxygen that unite with unit weight of sulphur. 

12. The two chlorides of tin have the following composition: The 
first chloride of tin contains 62.63% of tin and 37.37% of chlorine; 
the second chloride of tin contains 45.59% of tin and 54.41% of chlo- 



FIRST YEAR CHEMISTRY 365 

rine. Calculate the simple ratio between the varying amounts of chlo- 
rine that unite with unit weight of tin. 

13. The two sulphides of tin have the following composition: Stan- 
nous sulphide contains 78.80% of tin and 21.20% of sulphur; stannic 
sulphide contains 65.02% of tin and 34.98% of sulphur. Calculate 
the simple ratio between the varying amounts of sulphur that unite 
with unit weight of tin. 

14. The three oxides of chromium have the following composition: 
The first oxide of chromium contains 68.42% of chromium and 31.58% 
of oxygen; the second oxide of chromium contains 61.90% of chro- 
mium and 38.10% of oxygen; the third oxide of chromium contains 
52.00% of chromium and 48.00% of oxygen. Calculate the simple 
ratio between the varying amounts of oxygen that unite with unit 
weight of chromium. 

15. The three oxides of cobalt have the following composition: 
The first oxide of cobalt contains 78.66% of cobalt and 21.34% of 
oxygen; the second oxide of cobalt contains 73.44% of cobalt and 
26.56% of oxygen; the third oxide of cobalt contains 71.08% of co- 
balt and 28.92% of oxygen. Calculate the simple ratio between 
the varying amounts of oxygen that unite with unit weight of cobalt. 

16. The three oxides of copper have the following composition: 
The first oxide of copper contains 88.81% of copper and 11.19% of 
oxygen; the second oxide of copper contains 79.87% of copper and 
20.13% of oxygen; the third oxide of copper contains 66.49% of 
copper and 33.51% of oxygen. Calculate the simple ratio between 
the varying amounts of oxygen that unite with unit weight of copper. 

17. The three oxides of iron have the following composition: The 
first oxide of iron contains 77.77% of iron and 22.23% of oxygen; the 
second oxide of iron contains 72.41% of iron and 27.59% of oxygen; 
the third oxide of iron contains 70.00% of iron and 30.00% of oxy- 
gen. Calculate the simple ratio between the varying amounts of 
oxygen that unite with unit weight of iron. 

18. The three sulphides of iron have the following composition: 
The first sulphide of iron contains 63.63% of iron and 36.37% of sul- 
phur; the second sulphide of iron contains 53.84% of iron and 46.16% 
of sulphur; the third sulphide of iron contains 46.66% of iron and 
53.37% of sulphur. Calculate the simple ratio between the varying 
amounts of sulphur that unite with unit weight of iron. 

19. The three oxides of manganese have the following composi- 
tion: The first oxide of manganese contains 77.5% of manganese 
and 22.5% of oxygen; the second oxide of manganese contains 69.6% 



366 FIRST YEAR CHEMISTRY 

of manganese and 30.4% of oxygen; the third oxide of manganese 
contains 63.2% of manganese and 36.8% of oxygen. Calculate the 
simple ratio between the varying amounts of oxygen that unite with 
unit weight of manganese. 

20. The three oxides of nickel have the following composition; 
The first oxide of nickel contains 78.37% of nickel and 21.63% of 
oxygen; the second oxide of nickel contains 73.10% of nickel and 
26.90% of oxygen; the third oxide of nickel contains 70.73% of nickel 
and 29.27% of oxygen. Calculate the simple ratio between the vary- 
ing amounts of oxygen that unite with unit weight of nickel. 

21. The three oxides of tin have the following composition: The 
first oxide contains 88.14% of tin and 11.86% of oxygen; the second 
oxide of tin contains 83.21% of tin and 16.79% of oxygen, the third 
oxide of tin contains 78.81% of tin and 21.19% of oxygen. Calcu- 
late the simple ratio between the varying amounts of oxygen that 
unite with unit weight of tin. 

22. The five oxides of lead have the following composition: The 
first oxide of lead contains 96.27% of lead and 3.73% of oxygen; 
the second oxide of lead contains 92.82% of lead and 7.18% of oxy- 
gen; the third oxide of lead contains 90.65% of lead and 9.35% of 
oxygen; the fourth oxide of lead contains 89.61% of lead and 10.39% 
of oxygen; the fifth oxide of lead contains 86.61% of lead and 13.39% 
of oxygen. Calculate the simple ratio between the varying amounts 
of oxygen that unite with unit weight of lead. 

Problems Involving the Law of Boyle 

1. A gas measures 1546 c.c. at 754.6 mm. pressure. Find its vol- 
ume at 760 mm. pressure. 

2. A gas measures 650 c.c. at 765 mm. pressure. Find its vol- 
ume at 760 mm. pressure. 

3. ,A gas measures 1980 c.c. at 752 mm. pressure. Find its vol- 
ume at 774.2 mm. pressure. 

4. A gas measures 142 c.c. at 769.4 mm. pressure. Find its vol- 
ume at 749.3 mm. pressure. 

5. A gas measures 1782 c.c. at 760 mm. pressure. Find its vol- 
ume at 772 mm. pressure. 

6. A gas measures 980 c.c. at 760 mm. pressure. Find its vol- 
ume at 749.2 mm. pressure. 

7. A gas measures 480 c.c. at 75.1 cm. pressure. Find its volume 
at 760 mm. pressure. 






FIRST YEAR CHEMISTRY 367 

S. A gas measures 600 c.c. at 76.9 cm. pressure. Find its volume 
at 760 mm. pressure. 

9. A gas measures 654 c.c. at 74.9 cm. pressure. Find its volume 
at 76.5 cm. pressure. 

10. A gas measures 321 c.c. at 76.81 cm. pressure. Find its vol- 
ume at 750 mm. pressure. 

11. .A gas measures 492 c.c. at 760 mm. pressure. Find its volume 
at 77.5 cm. pressure. 

12. A gas measures 763.4 c.c. at 76 cm. pressure. Find its vol- 
ume at 755 mm. pressure. 

13. A gas measures 0.84 liters at 750 mm. pressure. Find its vol- 
ume at 760 mm. pressure. 

14. A gas measures 3.2 liters at 771.4 mm. pressure. Find its 
volume at 760 mm. pressure. 

15. A gas measures 1.7 liters at 759.3 mm. pressure. Find its 
volume at 771.2 mm. pressure. 

16. A gas measures 2.4 liters at 7S1 mm. pressure. Find its vol- 
ume at 751 mm. pressure. 

17. A gas measures 3.8 liters at 760 mm. pressure. Find its vol- 
ume at 781 mm. pressure. 

18. A sras measures 2.2 liters at 760 mm. pressure. Find its vol- 
ume at 749.9 mm. pressure. 

19. The air inclosed in the short arm of the Boyle tube measures 
16.6 c.c. when the barometer stands at 755 mm. Calculate the vol- 
ume of the inclosed air (a) at 760 mm., (b) at 770 mm., (c) at 750 mm., 
(d) at 2.5 atmospheres, when each atmosphere is 755 mm., and (e) the 
volume at 2.5 atmospheres when one atmosphere is 760 mm. 

20. The air inclosed in the short arm of the Boyle tube measures 
IS. 6 c.c. when the pressure is 761 mm. Calculate the volume of the 
inclosed air (a) at 760 mm., (b) at 750 mm., (c) at 770 mm., (d) at 
2.5 atmospheres of 761 mm. each, and (e) at 2.5 atmospheres of 760 mm. 
each. 

22. The volume of air inclosed in the short arm of the Boyle tube 
at a pressure of 715.5 mm. is 19.2 c.c. Find the volume of the in- 
closed air (a) at 766 mm., (b) at 2.4 atmospheres of 754.5 mm. each, 
and (c) at 3.6 atmospheres of 760 mm. each. 

Problems Involving the Law of Dalton 

1. A pas measures 100 c.c. at 7°C. Find its volume at 0°C. 

2. A gas meae iree 575 c.c. ;it 18°C. Find it- volume ;it 0°C. 

3. A gas measures 650 c.c. at 20°C. Find its volume at 0°C. 



368 FIRST YEAR CHEMISTRY 

4. A gas measures 600 c.c. at 9°C. Find its volume at 7°C. 

5. A gas measures 300 c.c. at 0°C. Find its volume at 17°C. 

6. A gas measures 450 c.c. at 0°C. Find its volume at 17°C. 

7. A gas measures 700 c.c. at 20°C. Find its volume at 30°C. 

8. A gas measures 1200 c.c. at 37°C. Find its volume at 27°C. 

9. A gas measures 250 c.c. at 20°C. Find its volume at 15°C. 

10. A gas measures 1000 c.c. at — 23°C. Find its volume at 0°C. 

11. A gas measures 800 c.c. at — 33°C. Find its volume at 0°C. 

12. A gas measures 264 c.c. at — 5°C. Find its volume at 0°C. 

13. In verifying the Law of Dalton, a student obtained the fol- 
lowing data: The weight of flask and fittings empty equalled 125.7 g.; 
the weight of flask and fittings with the water that ran in equalled 
215.3 g.; and the weight of flask and fittings full of water equalled 
465.2 g. From the above data, calculate (a) the volume of air at 
100°C, (b) the contraction of the air in going from 100° to 0°, (c) the 
volume of air at 0°C, (d) the amount the air expanded in going from 
0° to 100°, (e) the amount the air expanded in going from 0° to 1°, 
and (f) the amount 1 c.c. of air would expand in going from 0° to 1°. 
Carry the last answer out to five decimal places, but thruout the 
work save only one decimal place. 

14. In verifying the Law of Dalton, a student obtained the fol- 
lowing data: The weight of flask and fittings empty equalled 80.5 g.; 
the weight of flask and fittings after water has run in equalled 156.9 g.; 
and the weight of flask and fittings full of water equalled 372.4 g. 
From the above data, calculate (a) the volume of air at 100°C, (b) the 
contraction of the air in going from 100°C. to 0°C, (c) the volume 
of air at 0°C, (d) the amount the air expanded in going from 0° to 
100°, (e) the amount the air expanded in going from 0° to 1°, and 
(f) the amount 1 c.c. of air would expand in going from 0° to 1°. Carry 
the last answer out to 5 decimal places, but thruout the work save 
only 1 decimal place. 

Problems Involving both the Law of Boyle and the Law of Dalton 

1. A gas measures 637 c.c. at 755 mm. pressure and 17°C. Find 
its volume at 760 mm. pressure and 0°C. 

2. A gas measures 700 c.c. at 760 mm. pressure and 0°C. Find 
its volume at 770 mm. pressure and 17°C. 

3. A gas measures 800 c.c. at 750 mm. pressure and 17°C. Find 
its volume at 765 mm. pressure and 27°C. 

4. A gas measures 100 liters at 14°C. and 750 mm. pressure. Find 
its volume at 760 mm. pressure and 0°C. 



FIRST YEAR CHEMISTRY 369 

Problems on the Weight and Specific Gravity of Gases 

1. In determining the weight and specific gravity of hydrogen, 
a student obtained the following data: Weight of flask and fittings 
full of air = 91.6 g. Weight of flask and fittings full of water = 407.0 g. 
Barometer reading = 75S.0 mm., and thermometer reading = 22.8°C. 
The loss of weight due to hydrogen = 0.35 g. 1 c.c. of air at 0°C. 
and 760 mm. pressure weighs 0.00129 g. Calculate the specific gravity 
of hydrogen referred to air. Calculate also the weight of 1 c.c. of 
hydrogen at 0°C. and 760 mm. pressure. 

2. In determining the weight and specific gravity of illuminating 
gas. a student obtained the following data: Weight of flask and fit- 
tings full of air = 91.6 g. Weight of flask and fittings full of water 
= 407.0 g. Barometer reading = 758.0 mm. Thermometer read- 
ing = 22.8°C. The loss in weight due to illuminating gas = 0.20 g. 1 c.c. 
of air at 0°C. and 760 mm. pressure weighs 0.00129 g. Calculate the 
specific gravity of illuminating gas referred to air. Calculate 
the weight of 1 c.c. of illuminating gas at 0°C. and 760 mm. pressure. 

3. In determining the weight and specific gravity of chlorine, a 
student obtained the following data: Weight of flask and fittings 
full of air = 103.4 g. Weight of flask and fittings full of water = 
424.9 g. Barometer reading = 777.8 mm. Thermometer reading 
= 18.0°C. The gain in weight due to chlorine = 0.6 g. 1 c.c. of 
air at 0°C. and 760 mm. pressure weighs 0.00129 g. Calculate the 
specific gravity of chlorine referred to air. Calculate also the specific 
gravity of chlorine referred to hydrogen, assuming that air is 14.37 
times as heavy as hydrogen. 

4. In determining the weight and , specific gravity of bromine 
vapor, a student obtained the following data: Weight of flask and 
fittings full of air = 126.0 g. Weight of flask and fittings full of 
water = 476.0 g. Barometer reading = 770 mm. Thermometer 
reading = 27°C. The gain in weight due to bromine vapor = 1.885 g. 
1 c.c. of air at 0°C. and 760 mm. pressure weighs 0.00129 g. Calcu- 
late the specific gravity of bromine vapor referred to air. Calculate 
also the specific gravity of bromine vapor referred to hydrogen, as- 
suming that air is 14.37 times as heavy as hydrogen. 

Combining Numbers 

1. In determining the combining number for zinc, the following 
data were obtained: Weight of zinc = 1.95 g. The volume of hy- 
drogen evolved = 747.0 c.c. The barometer reading = 769.3 mm 



370 FIRST YEAR CHEMISTRY 

The thermometer reading =• 27°C. The correction for the pressure 
of aqueous vapor at 27°C. = 27 mm. The weight of one cubic centi- 
meter of hydrogen at 0°C. and 760 mm. pressure = 0.00009 g. Cal- 
culate the combining number of zinc and carry the answer out to one 
decimal place, saving one decimal place thruout the work on the re- 
duction of the gas volumes to S. T. P., and as many decimal places 
in the rest of the work as you think consistent with the above. 

2. In determining the combining number for zinc, the following 
data were obtained: Weight of zinc used = 1.30 g. The volume 
of hydrogen evolved = 498.0 c.c. The barometer reading = 769.3 
mm., and the thermometer reading = 27°C. The correction for the 
pressure of aqueous vapor at 27°C. = 27 mm. The weight of 1 c.c. 
of hydrogen at 0°C. and 760 mm. pressure = 0.00009 g. Calculate 
the combining number of zinc, and carry the answer out to one deci- 
mal place, saving one decimal place thruout the work on the reduc- 
tion of the gas volume to S. T. P., and as many decimal places in the 
rest of the work as you think consistent with the above. 

3. By putting 0.575 g. of sodium in water, 277.7 c.c. of hydrogen 
are obtained if the conditions are 0°C. and 760 mm. pressure. As- 
suming that no correction for the pressure of aqueous vapor is neces- 
sary, calculate the combining number for sodium. 1 c.c. of hydro- 
gen at 0°C. and 760 mm. pressure weighs 0.00009 g. 

4. 14.0 g. of iron evolve 0.5 g. of hydrogen. Find the combining 
number of iron. Also find the atomic weight of iron, using the Law 
of Dulong and Petit. The specific heat of iron = 0.1138. 

5. 9.0 g. of aluminium evolve just 1 g. of hydrogen. Find the 
combining number for aluminium; also find the atomic weight of 
aluminium, assuming that the valence of aluminium is 3. 

6. 3.95 g. of copper when heated to complete oxidation take on just 1 
g. of oxygen. The combining number of oxygen referred to hydrogen 
is 8. Find the combining number of copper referred to hydrogen. 

7. In mercury oxide 50.0 g. of mercury are combined with 4.0 of 
oxygen. Find the combining number of mercury referred to oxygen. 
Assuming that the combining number of oxygen referred to hydro- 
gen is 8, calculate the combining number of mercury referred to hy- 
drogen; also find the atomic weight of mercury, assuming that the 
valence of mercury is 2. 

8. A quantitative analysis of one of the oxides of lead showed that 
50 g. of the oxide contained 3.8647 g. of oxygen. Assuming that the 
atomic weight of oxygen is 16, and that the oxygen and the lead 
are united atom to atom, calculate the atomic weight of lead. 



FIRST YEAR CHEMISTRY 371 

Determination of Molecular Weights by the Physical Method 

The figures given below indicate the specific gravities of the sub" 
stances referred to hydrogen. From these specific gravities calculate 
the molecular weights of the different substances: 

Ammonia gas 8.5 Mercury vapor 100 

Argon 20 Sulphur dioxide 32 

Arsenic hydride 39 Sulphur trioxide 40 

Bromine vapor 79 Carbon bisulphide 38 

Hydrobromic acid gas 40.5 Nitrogen 14 

Hydrocyanic acid gas 13.55 Nitrous oxide 22 

Fluorine 19 Nitric oxide 15 

Hydrofluoric acid gas 10 Nitrogen dioxide 23 

Iodine vapor 127 Hydrogen telluride 64.5 

Hydriodic acid 64 Hydrogen selenide 40.55 

Carbon monoxide 14 Hydrogen sulphide 17 

Carbon dioxide 22 Marsh gas 8 

Ozone 24 Olefiant gas 14 

Phosphorus hydride 17 Acetylene 13 

Hydrochloric acid gas 18.25 Chlorine 35.5 

Determination of Molecular Weights by the Chemical Method 

1. 2.0 g. of potassium chlorate when heated till all the oxygen 
was driven off, lost 0.79 g. in weight. Assuming that each molecule 
of potassium chlorate loses three atoms of oxygen, and knowing that 
the atomic weight of oxygen is 16, calculate the molecular weight 
of potassium chlorate. Calculate also the molecular weight of po- 
tassium chloride that remained. Using the molecular weight of 
potassium chloride just determined, and knowing that the atomic weight 
of chlorine is 35.5, calculate the atomic weight of potassium chloride. 

2. If 1.20 g. of potassium chloride are treated with sulphuric acid 
till all the potassium chloride is transformed to potassium sul- 
phate, and the excess of sulphuric acid is all driven off, 1.42 g. of po- 
tassium sulphate remain. Assuming that the molecular weight of 
potassium chloride is 73.5 and that two molecules of potassium chlo- 
ride are needed to form one molecule of potassium sulphate, calcu- 
late the molecular weight of potassium sulphate. 

3. If dry hydrogen gas is passed over 15.9 g. of copper oxide in 
a hard jrlass tube till the copper oxide is completely reduced to cop- 
per, the resulting copper is found to weigh 12.7 g. Knowing that 



372 FIRST YEAR CHEMISTRY 

the atomic weight of oxygen is 16, calculate the molecular weight 
of copper oxide; also calculate the atomic weight of copper. 

4. If 12.5 g. of pure calcium carbonate are heated till all the car- 
bon dioxide has been driven off, there is a loss in weight of 5.5 g. Cal- 
culate the molecular weight of calcium carbonate. Calculate also 
the molecular weight of calcium oxide. Calculate also the atomic weight 
of calcium. The specific gravity of carbon dioxide referred to hydrogen 
is 22, and the specific gravity of oxygen referred to hydrogen is 16. 

5. 10 g. of the chloride of a certain univalent element are found 
to contain 60.6 per cent of chlorine. If the atomic weight of chlorine 
is 35.5 what is the atomic weight of the other element? 

Problems in Plain Stoichiometry 

1. How many grams of potassium sulphate will be necessary to 
react with 100 g. of barium nitrate to produce barium sulphate ? How 
many grams of barium sulphate will be precipitated? How many 
grams of potassium nitrate will be formed? Carry each answer out 
to two decimal places, and check the accuracy of your work by ap- 
plying the Law of Conservation of Mass. Ba = 137 N = 14 O = 16 
S=32 K=39 

2. How many grams of sodium chloride will 100 g. of sodium hy- 
drate give by neutralization with hydrochloric acid? Na=23 = 16 
H = l CI =35.5 

3. How many grams of potassium chloride will 100 g. of potassium 
hydrate give by neutralization with hydrochloric acid ? K = 39 O = 16 
H = l CI =35.5 

4. How many grams of ammonia gas, NH 3 , can be obtained from 
10 g. of ammonium sulphate by treatment with calcium hydrate? 
N=14 H = l S=32 = 16 Ca = 40 

5. If 100 g. of calcium fluoride are heated with sulphuric acid, 
how many grams of hydrofluoric acid can be formed? How many 
grams of sulphuric acid will be needed for the transformation? How 
many grams of calcium sulphate will be formed as a side product? 
Carry each answer out to two decimal places and check the accuracy 
of your work by applying the Law of Conservation of Mass. Ca = 40 
F = 19 H = l S=32 = 16 

Problems in Stoichiometry Involving Gas Volumes 

1. How many liters of carbon dioxide at 760 mm. pressure and 
17°C. may be obtained from 100 g. of calcium carbonate by treat- 



FIRST YEAR CHEMISTRY 373 

ment with hydrochloric acid? One liter of carbon dioxide at 0°C. 
and 760 mm. pressure weighs 1.9 g. Ca=40 C = 12 = 16 H = l 
CI =35.5 

2. How many liters of hydrochloric acid gas may be obtained from 
500 g. of sodium chloride at 770 mm. pressure and 17°C. by treatment 
with sulphuric acid? One liter of hydrochloric acid gas at 0°C. and 
760 mm. pressure weighs 1.63 g. Na=23 CI =35.5 H = l S=32 O = 16 

3 How many liters of hydrogen gas measured at 15°C. and 765 mm. 
pressure can be obtained by treating 200 g. of metallic zinc with 
dilute sulphuric acid? One liter of hydrogen gas at 0°C. and 760 
mm. pressure weighs 0.09 g. H = l S=32 = 16 Zn=65.5 

4. How many c.c. of hydrogen sulphide gas measured at 14°C. 
and 750 mm. pressure may be obtained from 50 g. of iron sulphide 
by treatment with hydrochloric acid? One liter of hydrogen gas at 
0°C. and 760 mm. pressure weighs 0.09 g., and the specific gravity 
of hydrogen sulphide gas referred to hydrogen is 17.0. Fe =56 S =32 
H = l CI =35.5 

Determination of Chemical Formulae from Percentage Composition 

1. Ether has the following composition: Carbon 65%, hydrogen 
14%, and oxygen 21%. One liter of ether vapor at 0°C. and 760 mm. 
pressure weighs 3.33 g. One liter of hydrogen gas at the same con- 
ditions weighs 0.09 g. Calculate the formula for ether. C = 12 H = 1 
= 16 

2. Photographers' 'hypo has the following composition: So- 
dium 29%, sulphur 40.5%, and oxygen 30.5%. Molecular weight of 
hypo is 158. Calculate its formula. Na=23 S=32 = 16 

3. A compound composed of carbon and hydrogen has the fol- 
lowing composition: Carbon 80%, hydrogen 20%. The vapor den- 
sity of the compound is 15. Calculate the molecular formula of the 
compound. C = 12 H = l 

4. A compound contains 92.3% of carbon and 7.7% of hydrogen. 
Its vapor density is 2.44 times as heavy as oxygen under the same 
conditions, and oxygen is 16 times as heavy as hydrogen. What 
is the formula of the compound? H = l C = 12 = 16 

Determination of the Percentage Composition of Compounds from 
their Molecular Formulae 

1. Calculate the per ceni of water of crystallization in crystal- 
lized barium chloride, BaCl 2 .2H 2 0. Ba = 137 CI =35.5 H = l = 16 



374 FIRST YEAR CHEMISTRY 

2. Calculate the per cent of S0 2 in anhydrous sodium sulphite, 
Na 2 S0 3 . Na=23 S = 32 = 16 

3. The formula for ordinary alum is K 2 S0 4 .A1 2 (S0 4 ) 3 .24H 2 0. Cal- 
culate the per cent of water of crystallization in alum. Also calcu- 
late the per cent of metallic aluminium in alum. Also calculate 
the per cent of metallic potassium in alum. Also calculate the per 
cent of S0 3 in alum. K=39 S=32 = 16 Al=27 H = l Carry 
each quotient out to four decimal places and then express the answer, 
using the per cent sign. 

4. Calculate the percentage composition of marsh gas, CH 4 . 

5. Calculate the percentage composition of acetylene, C 2 H 2 . 

6. Calculate the percentage composition of acetic acid, C 2 H 4 2 . 

7. Calculate the percentage composition of cane sugar, Ci 2 H 22 0n. 

8. Calculate the percentage composition of galena, PbS. 

9. Glauber's Salt has the formula Na 2 SO 4 .10H 2 O. Calculate the 
per cent of water of crystallization in this compound. Also calcu- 
late the per cent of metallic sodium; also calculate the per cent 
of S0 3 . Also calculate what per cent of the crystallized salt is an- 
hydrous sodium sulphate. Na = 23 S = 32 = 16 H = 1 



Practical Questions 

The following questions are designed to make a student 
think independently, and to use the knowledge that he 
has gained from the use of this book in the laboratory. 

1. Is every case of oxidation accompanied by heat? Is every 
case of oxidation accompanied by flame? 

2. Is it possible to oxidize wood ashes? Is it possible to oxidize 
coal ashes? 

3. Why does alcohol give little or no smoke when it burns? 

4. Why does tar produce much smoke when it burns ? 

5. Why does blowing on a wood fire increase the combustion? 
Why does blowing on a candle extinguish the flame ? 

6. Why does a Bunsen burner flame give great heat, little light, 
and no smoke? 

7. Why is the flame of a Bunsen burner hotter when the holes at 
the bottom are open than when they are closed? 

8. Why does soft wood ignite more easily than hard wood? 






FIRST YEAR CHEMISTRY 375 

9. Why is galvanized iron rather than plain sheet iron used in 
conductors on houses? 

10. If several pounds of metallic sodium caught fire in a wooden 
building, would you try to extinguish the fire by pouring water on 
the burning sodium? What would be a better extinguisher for this 
kind of a fire ? 

11. What is the film which gathers on the chimney of a lamp 
when the lamp is first lighted? Why does this film soon disappear? 

12. Name one example of oxidation occurring in the body. 

13. What would happen if the atmosphere consisted of undiluted 
oxygen gas ? 

14. What effect would an atmosphere of nitrogen have upon a 
human being? 

15. What compound is removed from a room by means of ventilation ? 

16. What burns in the case of a "chimney on fire"? 

17. Assuming that carbon dioxide causes the foulness of air, and 
remembering that carbon dioxide is heavier than air, why is it that 
a room can best be ventilated by opening a window a little at the 
top and a little at the bottom? 

18. Ought the doors of a burning house to be thrown open? 

19. How can you prove that hydrogen is lighter than air if the only 
apparatus you have on hand consists of two test tubes, some matches, 
and a rubber bag full of hydrogen? 

20. How can you prove that carbon dioxide is heavier than air, 
if the only apparatus you have on hand consists of two test tubes, 
some matches, and a bag of carbon dioxide? 

21. How could you show whether an old well was full of carbon 
dioxide? 

22. How can you prove that a candle flame is hollow? 

23. Suppose that you have three unlabeled fruit jars, one con- 
taining oxygen, another containing ammonia gas, and the third con- 
taining hydrochloric acid gas. What is the easiest test to apply in 
order to determine which is which? 

24. Why would not pure hydrogen make a good substitute for 
illuminating gas? Give two reasons. 

2o. What metals are mentioned in the Bible? What alloys are 
mentioned in the Bible? 

26. Is an amaljram an alloy ? Is every alloy an amalgam? 

27. What metal is liquid ;it ordinary temperature? 

28. Why does silver money turn dark when carried in the same 
pocket with sulphur match' »? 



376 FIRST YEAR CHEMISTRY 

29. Mention four common metals that will float on mercury. 

30. Why can we not write on the black board with marble as well 
as with chalk, tho both are calcium carbonate? 

31. How may lime-water be made from oyster shells? 

32. If you ask for "potash" at a drug store, what compounds 
of potassium might you get 

33. If a dime be dissolved in nitric acid, the solution is blue. Why 
is this so? 

34. Why does ordinary white paint turn black more quickly in 
a chemical laboratory than in a physical laboratory ? 

35. In generating carbon dioxide from marble, why is it better 
to use hydrochloric acid than sulphuric acid ? 

36. Why must carbonic acid be kept in an air-tight bottle ? 

37. Why are the sparks between the trolley wheel of an electric 
car and the overhead wire often green? 

38. Why should a bottle of concentrated aqua ammonia not be 
kept near a hot radiator? 

39. Why does aqua ammonia lose its strength on exposure to air? 

40. What causes the presence of dissolved salts in the ocean? 

41. Which contains the greater amount of metallic sodium, a 
kilogram of crystallized sodium carbonate or a kilogram of anhydrous 
sodium carbonate? 

42. Why must sodium amalgam not be exposed to air when not 
in use? 

43. How may water suitable to drink be made from sea- water? 

44. Why must a jar containing stick sodium hydrate be kept 
tightly stoppered when not in use? 

45. Why must all the air be expelled from a hydrogen generator 
before the gas is lighted at the end of the delivery tube? What is 
the best way to light the hydrogen ? 

46. Why do we not consider water an element in spite of the fact 
that the alchemists so considered it? 

47. Why does rock candy not taste so sweet as a syrup made by 
dissolving rock candy in water? 

48. How may silver nitrate stains be removed from the skin ? 

49. Why are bean pots glazed? Why are flower pots not glazed? 

50. Why do eggs tarnish silver spoons? 

51. What is the best remedy for acid stains on cloth? Why must 
this remedy be applied while the stains are fresh? 

52. Why is distilled water insipid to the taste? 

53. Why is a lead pencil called a lead pencil in spite of the fact 



FIRST YEAR CHEMISTRY 377 

that there is neither metallic lead nor any lead compound in its com- 
position? 

54. Must an acid always be sour? Name one that is not sour. 

List of Text-Books for Reference Library 

The following text-books are recommended for the use 
of the teacher with this text. None of the numerous ele- 
mentary texts are included in this list. 

Richter, Text-book of Inorganic Chemistry. Edited by Klinger. 
Translated by Smith. 430 pages. P. Blakiston's Son and Co. Splen- 
did reference book on General Descriptive Inorganic Chemistry. 
Theory treated briefly but clearly. 

Remscn. A College Text-book of Chemistry. 690 pages. Henry 
Holt and Co. Treats of both theory and descriptive chemistry. 

.4. Smith. Introduction to General Inorganic Chemistry. 780 pages. 
The Cent urs* Co. Very full. Lays special emphasis on definitions. 

Newell, Descriptive Chemistry. 590 pages. D. C. Heath and Co. 
Very full. Interesting. 

Storer and Lindsay, An Elementary Manual of Chemistry. 450 
pages. American Book Co. 90 pages are devoted to a broad and 
interesting treatment of organic chemistry. 

F. H. Thorp, Outlines of Industrial Chemistry. 540 pages. The 
Macmillan Co. Illustrated. Comprehensive accounts of the more 
important industries. A splendid reference work. 

V enable, A Short History of Chemistry. 160 pages. D. C. Heath 
and Co. Best short treatment of the history of the subject. 

Comey, A Dictionary of Chemical Solubilities (Inorganic). 515 
pages. The Macmillan Co. Very full. Until recently practically 
the only reference work on solubilities outside of the brief tables found 
in some text-books on qualitative analysis. 

Biedermann, Chemiker-Kalendar. 362 pages. Julius Springer, 
of Berlin. Published annually. Contains many valuable tables. 
In German. 

Ostwald, Scientific Foundations of Analytical Chemistry. Trans- 
lated by McGowan. 215 pages. The Macmillan Co. Good treat- 
ment of modern theoretical conceptions of analytical reactions. 

Talbot and Blanchard, The Electrolytic Dissociation Theory and 
Some of its Applications. 84 pages. The Macmillan Co. An ele- 



378 FIRST YEAR CHEMISTRY 

mentary treatise for the use of students of chemistry. A valuable 
reference book for both student and teacher. 

Watts, Dictionary of Chemistry. Edited by Morley and Muir. 
4 volumes. Total number of pages 3290. Longmans, Green and Co. 
Described by the Text-book Committee of the New England Asso- 
ciation of Chemistry Teachers, as: "Voluminous, comprehensive, 
and authoritative. The best available reference book covering the 
whole subject. Probably will not be replaced for many years." 

T. E. Thorpe, A Dictionary of Applied Chemistry. 3 volumes. 
Total number of pages 1511. Longmans, Green and Co. It sup- 
plements Watt's Dictionary, treating of applied chemistry as that 
does of pure chemistry. 

Smith and Hall, The Teaching of Chemistry and Physics in the 
Secondary School. 375 pages. Longmans, Green and Co. Sug- 
gestive and helpful. Contains full lists of reference books. 

Journal of the American Chemical Society, The official organ of the 
American Chemical Society. Sent free to members of the American 
Chemical Society. 



List of Individual Apparatus 

List of individual apparatus. — Each student should be 
supplied with the following pieces of apparatus: 

Locker Key No 

Text-book. 

Laboratory notebook. See page 1 for description. The author 
finds it advisable to have the notebooks for his classes made to 
order by Carter, Rice and Co. of Boston. 

Ring stand and three rings. Stand 24 inches high with 5 by 7 base; 
the rings 1 inch, 2 inches, and 3 inches in diameter respectively. 

Two clamps for stand. 

Tripod. 

Four fruit jars and washers. "Lightning" make; 1 pint size. 

Three prescription bottles. 125 c.c, 250 c.c, and 500 c.c, respec- 
tively. 

Two catch bottles. 125 c.c. each. 

Nest of three beakers. 120, 240, and 360 c.c. are convenient sizes. 

Three flasks. 150, 250, and 500 c.c. 

Kjeldahl flask. 250 c.c; narrow neck. 



FIRST YEAR CHEMISTRY 379 

Graduated cylinder. 100 c.c. 

Two funnels. 3 inch and 4 inch. 

Thistle tube. 10 inch stem. 

Twelve test tubes. 4 inch, 5 inch and 6 inch; assorted 

One stick glass rod. 1 meter. 

Two sticks glass tubing. 1 meter each. 

Mortar and pestle. 3.5 inches diameter. 

Porcelain evaporating dish. 8 cm. diameter. 

Porcelain crucible and cover. 1 1-2 inch diameter 

Nest of four Hessian crucibles. 

Test tube rack. 

Test tube brush. 

Bunsen burner and hose. 

Bat-wing burner and hose. 

Box of weights. See page 11 for description. 

Twenty round filters. Ten 12.5 cm. diameter and ten 20 cm. diameter. 

Ten rubber connectors. 

Ten assorted corks. To fit different bottles and flasks. 

Round file. 5 inch. 

Triangular file. 5 inch. 

Iron rod. Piece of old telephone wire 20 cm. long. 

Brass forceps. 5 inch. 

Crucible tongs. 20 cm. 

Two iron gauzes. Coarse and fine wire netting cut in pieces 12 cm. 

square. 
Tubing clamp. See Fig. 1. 
Deflagrating spoon. See Fig. 1. 
Pipe-stem triangle. See Fig. 1. 
Horn spatula. 10 cm. long. 
Thirty centimeter rule. 
Two towels. 

List of General Apparatus 

The following apparatus should be on hand in addition 
to the individual apparatus given to each student. The 
number of pieces of each article in this list will depend, of 
course, upon the size of the class. 

Platform balance. See Fig. 4 
Iron weights. See Fig. 5. 



380 FIRST YEAR CHEMISTRY 

Horn-pan balance. See Fig. 6. 

Blast lamp. 

Richards blower or foot bellows. See Fig. 29 and Fig. 28. 

Pneumatic troughs. 

Wooden toothpicks. 

Electrolytic apparatus. See Fig. 45. 

Plunge battery or other source of electricity. 

Eudiometer and stand. See Fig. 49. 

Rhumkorff coil. 

Telegraphic key. 

Capillary glass tubing. 

Hard glass tubing, 7 mm. internal diameter. 

Hard glass tubing, 1 cm. internal diameter. 

Soft glass tubing, 4 mm. external diameter. 

Magnifying glass. 

Suction pump. 

Battery jars, large and small. 

Anvil and hammer. 

Asbestos sheet. 

6 by 1 test tubes. 

Rubber gas bags with brass stop- cocks. 

Labels. 

Retorts, 250 c.c. 

Meter rod. 

Boyle tube. See Fig. 74. 

High grade barometer. See Fig. 75. 

Copper boiler. 

Thermometer with Centigrade scale. 

Two-liter prescription bottle with one-hole rubber stopper to fit it. 

Heavy-walled rubber tubing. 

Gas balance. 

Bricks. 

Four liter prescription bottle with a two-hole rubber stopper to fit it. 

Kjeldahl flask with wide neck. 

Closed ignition tube. 6 inch. 

Platinum test wires. 

Sand baths. 

Pipettes, 10 c.c. capacity. 



FIRST YEAR CHEMISTRY 



381 



List of Chemicals 



List of chemicals.— The following is a complete list of 
all the chemicals needed for the work indicated in this book. 



alcohol 

alum 

ammonium carbonate 

ammonium chloride 

ammonium hydrate 

ammonium nitrate 

antimony 

arsenic 

arsenic oxide 

aluminium 

aluminium oxide 

aluminium sulphate 

barium chloride, crystallized 

barium nitrate 

bismuth 

bleaching powder 

bromine 

calcium 

calcium carbonate (marble) 

calcium chloride 

calcium fluoride 

calcium hydrate 

calcium oxide 

candles 

carbon (see Expt. 4.5) 

carbon dioxide, cylinder 

carbon bisulphide 

carbonic acid 

chlorine 

copper 

copper oxide 

copper sulphate, crystallized 

emery 

ether 

gypsum 

bydrobromic acid 



hydrochloric acid 

iodine 

iron 

iron chloride 

iron oxide, ferric 

iron oxide, ferroso- ferric 

iron sulphate, crystallized 

iron sulphide 

kerosene 

lead 

lead nitrate 

lead oxide, litharge 

lead oxide, red lead 

litmus, lumps 

litmus paper 

magnesium 

magnesium chloride 

magnesium oxide 

magnesium sulphate, crystallized 

manganese dioxide 

mercury 

mercury oxide 

nitric acid 

paraffine 

phorphorus, red and yellow 

plaster of Paris 

platinum sponge 

potassium 

potassium bromide 

potassium carbonate 

potassium chromate 

potassium chlorate 

potassium chlorate, pure 

potassium chloride 

potassium hydrate, stick 

potassium iodide 



382 FIRST YEAR CHEMISTRY 

potassium nitrate sodium sulphate, crystallized 

potassium sulphate, crystallized sulphur 

sand sulphuric acid 

silver tartar emetic 

silver nitrate tin 

soap, pure castile tin chloride 

sodium tin oxide 

sodium amalgam turmeric paper 

sodium carbonate, crystallized vaseline 

sodium chloride water, distilled 

sodium hydrate, stick zinc 

sodium peroxide zinc chloride 

sodium phosphate zinc oxide 

sodium silicate, syrup zinc sulphate, crystallized 

Suggestions to the Teacher 

The preface to this book sets forth the motif of the au- 
thor and indicates how the text is used in his classes. Many 
suggestions, therefore, that would naturally be looked for 
here will be found fully set forth in the preface. The fol- 
lowing suggestions are inserted, however, to emphasize 
certain points already touched upon, and to include hints 
which may be of help in using this text, but which are not 
given elsewhere. 

The instructor should always be within easy reach of 
the pupil for appeal in case of difficulty with text or ex- 
periments, but the pupil should be encouraged to solve 
his own difficulties as far as possible. 

The author finds that thirty-five school weeks are ample 
for doing the work laid out in this book. The number 
of recitations per week vary with the nature of the text. 
Four or five hours per week in the laboratory should en- 
able the student to cover the assigned work intelligently. 

Much valuable time, both of teacher and of pupil, may 
be saved by having everything on hand and ready for use 
when the classes come in to work. 






FIRST YEAR CHEMISTRY 383 

In the treatment of the notebooks, each teacher must 
be guided by the conditions existing in his own laboratory. 
The section on the laboratory notebook on page 1, ex- 
plains the method the author has found most satisfactory 
in his laboratory. 

It should be emphasized here that the notebook becomes 
the student's textbook. All essentials have been included 
in the text and the author feels that the directions are 
clear and full enough for a student of less than average 
ability. In spite of this, constant watchfulness on the 
part of the teacher is necessary to see that the student 
does not wander too far from the path laid out for him. 
Since the student learns his chemistry from his own ex- 
perience the use of other text-books is reduced to a minimum. 

The Chemical Investigation, together with the work 
that leads up to it, shows that the Atomic Theory is not 
absolutely necessary to an intelligent acquaintance with 
chemical substances and the changes they undergo. 

In his own classes the author has recently let the stu- 
dents read pages 183 to 197 on the Theory of Chemistry 
aloud without outside preparation. This enables the class 
to cover the ground more rapidly, and the general im- 
pression thus gained by the student helps to tie together 
the various parts of the theory which are treated in con- 
siderable experimental detail in pages 219 to 258. 

Is it necessary to say that the teacher should certainly 
try all the lecture experiments by himself before trying 
to perform them before the class? 

Orderliness, cleanliness, use of common sense, and will- 
ingness and ability to follow instructions faithfully and 
understandingly are essential to success in chemical work. 

It is also of the greatest importance that the student 
should agree with himself to see results exactly as they 
present themselves, and refuse to allow himself to be prej- 



384 FIRST YEAR CHEMISTRY 

udiced in his observation by what he may think is present, 
or by what he may think ought to happen. This attitude 
of absolute honesty is more easily attained by the student 
if he comes in contact with it in all tests that the instructor 
may find it necessary to make in connection with the stu- 
dent's work. 

A collection of the chemical elements and compounds 
as they occur in nature and as they are found in the labora- 
tory, is a distinct help to the students if exhibited in such, 
place and form as to be accessible to the members of the 
class. The author uses a collection of about fifteen hun- 
dred such specimens to good advantage. He has a lim- 
ited number of printed catalogs (74 page pamphlets) of 
this collection which he will gladly send to any teachers who 
are contemplating developing such a collection. The 
crystals illustrated in Figures 22, 23, 24, and 25 were se- 
lected from this collection. 

On the opening day each student is handed a printed 
slip containing the following General Directions: This slip, 
was prepared to fit the conditions in the author's labora- 
tory and is reprinted here simply to serve as a guide for 
other teachers who may wish to use a similar plan. 

The General Information contained on this slip is such as will ac- 
quaint the student with the running of the laboratory and enable 
him to get started in his work. 

The Bulletin Board is in the main laboratory, and on it are posted 
any necessary notices in regard to the work in chemistry. Students 
are held responsible for all notices posted on this bulletin board. 

The Storeroom furnishes all apparatus and chemicals needed in the 
work of the course; the text-book and notebook come with the ap- 
paratus obtained at the first laboratory period, unless these were dis- 
tributed at the first meeting of the class. The hours during which 
the storeroom is open are posted on the bulletin board. Each stu- 
dent should procure for himself a coat or apron to protect the clothes; 
this is the only article not furnished by the storeroom, 



FIRST YEAR CHEMISTRY 385 

The Lockers are assigned by lot, one to each student, and the list 
is posted on the bulletin board in the laboratory as soon as possible 
after the first meeting of the class. Each student should look on 
this list to ascertain the number of his locker; this number, together 
with his name and the date, he should put on the long printed slip 
containing the list of apparatus. This slip may be obtained from the 
instructor, and, when properly filled out, should be sent to the store- 
room by the elevator; the articles mentioned on the list will be re- 
ceived in return. Each locker has a closet and several drawers. If 
anything about the locker is out of order, report it at once in writing 
to the instructor. 

Look over the apparatus carefully, and familiarize yourself with 
the names of the different pieces of apparatus, referring to Fig. 1 
in the text-book if in doubt about the name of any piece; also 
see if the apparatus is in good condition. If any articles are missing 
or broken, report it to the instructor at once, Any mistakes made 
in preparing the sets of apparatus cannot be rectified if not reported 
within twenty-four hours after the apparatus has been delivered to 
the student. 

Care of apparatus and desk. — Put all things away in the locker 
and lock up. The key and the apparatus assigned to the stu- 
dent are in his keeping for the year. Articles returned in good 
condition at the end of the year are credited without charge; all 
apparatus not returned or not in good enough condition to 
be given other students, will be charged to the student who has 
ordered it. When the work for each day has been finished, all 
apparatus should be put away in the lockers, and the desk top 
should be left clear; any articles left out on the desks are removed 
to the storeroom when the building is inspected at the close of 
the day's work, and such articles can be regained only by ordering 
again. 

Chemicals. — All chemicals used during the year are put out on the 
shelves or in the end drawers as needed, and are for the general use 
of the class. When necessary the bottles may be taken to the in- 
dividual desks, but for the convenience of the next user they should 
be returned to the shelves as soon as possible. 

Order blanks. — These are to be used whenever any extra pieces 
of apparatus are needed. They should be properly filled out with 
the name of the student, the date, and the name of the article wanted, 
and then senl by the elevator to the storeroom. The articles wanted 
will be delivered immediately. 



386 FIRST YEAR CHEMISTRY 

Helps on the experiments. — The following helps on the 
experiments are for the use of the teacher, and treat of 
such details as the author has found essential for successful 
performance of the experiments. The numbers refer to 
the number of the experiment. 

5. Copper wire Nos. 18 and 26 are convenient sizes. 

11. With sensible care phosphorus experiments are not danger- 
ous. Work with phosphorus should under no circumstances be al- 
lowed at the individual desks, however, on account of. danger of fire 
from unused phosphorus residues. 

19. The instructor should see that the student observes all the 
details of this experiment. 

23. The author is indebted to Mr. W. E. Fiske, Instructor in Phys- 
ics at The Phillips Exeter Academy, for the use of his plunge batteries; 
these are of the carbon-zinc-sulphuric acid-chromic anhydride type 
and were made by E. W. Wescott, '07. 

24. If it should happen that the laboratory is not supplied with 
the apparatus necessary for Experiments 23 and 24, the instructor 
might omit the experimental part, but, in order to sustain the logical 
sequence of the experiments in this book, he should give the gist of 
Experiments 23 and 24 as a lecture and require the students to take 
notes thereon. The data obtained in these two experiments are re- 
ferred to in later experiments. 

25. The teacher should try this experiment himself before class 
and then watch the students closely that they connect the tubes at 
just the right moment. This experiment does not work well, if the 
apparatus is connected before generating steam, for the hard glass 
tube is likely to break before the magnesium gets hot enough. Fur- 
thermore, dry steam does not react readily with magnesium. 

32. 10 grams of zinc and 10 c.c. of acid are right proportions to 
get good crystals; therefore the zinc should be weighed and the acid 
should be measured. Burning hydrogen with the platinum tip is 
a good lecture experiment. 

37. Use No. 18 wire. 

42. Use fused globular iron sulphide. 

46. Even if the student is not able to get the properties of carbon 
dioxide from that which he makes in the jar, he can get them from 
the bag full of the gas, and this experiment will teach him the neces- 
sity for closeness of observation. The instructor might well try this 



FIRST YEAR CHEMISTRY 387 

experiment himself before class to see that the materials are in 
proper condition. See that the distilled water put out for this ex- 
periment has no air dissolved in it. If it has air in it. boil it out and 
then keep the cooled water in tightly stoppered bottles. The cyl- 
inders of carbon dioxide sold for soda fountain use are extremely con- 
venient sources of carbon dioxide for the laboratory. Carbonic acid 
may be obtained from the local druggist. It is well to have a supply 
of old lager beer bottles in which to keep this liquid. 

47. An interesting side experiment on carbon dioxide is to. make 
solid carbon dioxide and freeze mercury with it. 

48. The instructor would do well to try this experiment by him- 
self before class. 

49. Chlorine should be made up a day or two before it is needed; 
it may be made easily from bleaching powder and hydrochloric acid 
in the cold in an ordinaiy generator; it should be caught in dry pint 
fruit jars by displacement of air. For large classes make the generator 
of a large bottle and use a separatory funnel instead of a thistle tube. 

66. The sodium amalgam for the bottle on the shelf may be 
made by heating 500 grams of mercury in a Hessian crucible of about 
250 c.c. capacity till the mercury is at 200°C, dropping in a lump 
of sodium weighing 50 grams, and covering immediately with a heavy 
plate. The preparation of sodium amalgam needs great care; it 
should be done in a hood; the reaction is very vigorous. Pour the 
molten amalgam into a tin pan and, when cooled, break it into small 
pieces. Store it in air-tight jars. 

75. 100 gram lots of metallic calcium may be bought from Eimer 
and Amend in sealed glass tubes at about $2.00 per tube. Other 
dealers also probably sell it. 

79. Spread powdered quick lime out to contact with air till it 
becomes air slaked. 

98. In the author's laboratory the barometer is kept in the in- 
structor's office, but the thermometer is hung in the main laboratory 
near the office door where the temperature is nearly constant. 

121. Mossy lead is easy to make. Hold old lead pipe in the blast 
lamp flame and let the melted lead drop into water. 

143. An interesting modification of this experiment is to dissolve 
a dime in nitric acid and then precipitate the silver on sheet copper. 

155. The instructor should prepare the sodium silicate solution 
of such a strength that a few drops of concentrated hydrochloric acid 
causes a precipitate of silicic acid but does not fill the test tube with 
gelatinous acid. 



388 



FIRST YEAR CHEMISTRY 



Answers to Additional Problems 



Metric System 


6. 


1:2 


(b) 


16.27 c.c. 


1. 128.5 cm. 

2. 10.861 g. 

3. (a) 460 mm. 
(b) 4.6 dm. 


7. 

8. 

9. 
10. 
11. 


2:3 
1:2 
1:2 
1:2 
2:3 


(c) 
(d) 
(e) 
20. (a) 
(b) 


16.71 c.c. 

6.'64 c.c. 

6.59'c.c. 
17.48 c.c. 
17.71 c.c. 


(c) .46 m. 

(d) 18.4 in. 

(e) 1.533 ft. 


12. 
13. 
14. 


1:2 
1:2 
3:4:6 


(c) 
(d) 
(e) 


17.14 c.c. 
6.88 c.c. 
6.99 c.c. 


4. (a) 860 mm. 
(b) 8.6 dm. 


15. 
16. 


6:8:9 
1:2:4 


21.(a) 

(b) 


18.62 c.c. 
18.87 c.c. 


(c) .86 m. 

(d) 34.4 in. 

(e) 2.86 ft. 


17. 
18. 
19. 


6:8:9 
2:3:4 
2:3:4 


(c) 
(d) 
(e) 


18.38 c.c. 
7.44 c.c. 
7.44 c.c. 


5. 12000 c.c. 

6. (a) 625000 cu. mm. 

(b) .625 cu. dm. 

(c) .625 1. 


20. 
21. 
22. 


6:8:9 
2:3:4 
3:6:8:9:12 


22. (a) 
(b) 
(c) 


18.9 c.c. 
8 c.c. 
5.29 c.c. 


(d) .000625 cu. m. 
7. (a) .58 cu. dm. 




Law of Boyle 


Law 


of Dalton 


(b) .58 1. 


1. 


1535.0 c.c. 


1. 


97.5 c.c. 


8. 1.54 1. 


2. 


654.2 c.c. 


2. 


539.4 c.c. 


9. 139320 g. 


3. 


1923.2 c.c. 


3. 


605.6 c.c. 


10. (a) 450 cu. m. 


4. 


145.8 c.c. 


4. 


615.3 c.c. 


(b) 450,000 1. 


5. 


1754.3 c.c. 


5. 


318.6 c.c. 


(c) 44.775 cu. m. 


6. 


994.1 c.c. 


6. 


478.0 c.c. 


11. 800 g. 


7. 


474.3 c.c. 


7. 


723.8 c.c. 


12. (a) 240 g. water 


8. 


607.1 c.c. 


8. 


1161.2 c.c. 


(b) 300 g. glycerine 


9. 


640.3 c.c. 


9. 


245.7 c.c. 


(c) 432 g. H 2 S0 4 


10. 


328.7 c.c. 


10. 


1092.0 c.c. 




11. 


482.4 c.c. 


11. 


910.0 c.c. 


Law of Multiple 
Proportions 


12. 
13. 
14. 


768.5 c.c. 

828.9 c.c. 

3248.0 c.c. 


12. 
13. (a) 

(b) 


268.9 c.c. 

339.5 c.c. 

89.6 c.c. 


1. 3:5 


15. 


1673.7 c.c. 


(c) 


249.9 c.c. 


2. 3:5 


16. 


2495.8 c.c. 


(d) 


89.6 c.c. 


3. 3:5 


17. 


3697.7 c.c. 


(e) 


.896 c.c. 


4. 3:5 


18. 


2229.6 c.c. 


(0 


.00358 c.c. 


5. 1:2 


19.( 


a) 16.49 c.c. 


14. (a) 


291.9 c.c. 



FIRST YEAR CHEiMISTRY 



389 



(b) 76.4 c.c. 

(c) 215.5 c.c. 

(d) 76.4 c.c. 

(e) .764 c.c. 

(f) .003545 c.c. 

Laws of Boyle and 
Dalton 

1. 595.7 c.c. 



Molecular Weights by 
Physical Method 



733.9 c.c. 
811.3 c.c. 
93.79 I. 



Weight and Specific 
Gravity of Gases 

1. .06503 
0.0000839 g. 

2. 0.46 
0.00060 g. 

3. 2.48 
35.64 

4. 5.52 
79.32 

Combining Numbers 



1. 
2. 
3. 
4.(a) 

(b) 
5. (a) 

(b) 
6. 
7.fa) 

(b) HK) 

(c) 200 
207.0 



32.6 
32.6 
23.0 
28 
56 
9 
27 
31.6 
12.5 



NH 3 

A 

AsH 3 

Br 2 

HBr 

HCN 

F 2 
HF 

I2 
HI 
CO 
C0 2 

O3 

PH 3 

HC1 

Hg 

S0 2 

S0 3 

CS 2 

N 2 

N 2 

NO 

N0 2 

H 2 Te 

H 2 Se 

H 2 S 

CH 4 

C 2 H 4 

C 2 H 2 

Cl 2 



17 

40 

78 
158 

81 

27.1 

38 

20 
254 
128 

28 

44 

48 

34 

36.5 
200 

64 

80 

76 

28 

44 

30 

46 
129 

81.1 

34 

16 

28 

26 

71 



Molecular Weights by 
Chemical Method 

l.(a) 121.5 

(b) 73.5 

(c) 38 

2. 173.95 



3. (a) 
(b) 

4. (a) 
(b) 
(e) 

5. 



79.5 
63.5 
100 
56 
40 
23 



Plain Stoichiometry 



BaS0 4 
K 2 S0 4 
KNO3 



CaS0 4 

HF 

H 2 S0 4 



89.27 g. 
66.66 g. 
77.39 g. 
146.25 g. 
133.03 g. 
2.5 g. 
174.36 g. 

51.28 g. 
125.64 g. 



Stoichiometry Involving 
Gas Volumes 

1. 24.53 I. 

2. 200.6 1. 

3. 70.89 1. 

4. .013407 c.c. 

Chemical Formulae from 
Percentage Composition 

1. C 4 H 10 O. 

2. Na 2 S 2 3 

3. C 2 H 6 

4. C 6 H 6 

Percentage Composition 
from Molecular Formula 

1. 14.75% 

2. 50.7% 

3. H 2 0, 45.56% 
A!, 5.69% 



390 



FIRST YEAR CHEMISTRY 



K, 8.22% 


6. C, 40% 


8. Pb, 86.6% 


S0 3 , 8.43% 


H, 6.66% 


S, 13.4% 


C, 75% 


0, 53.33% 


9. H 2 0, 55.90% 


H, 25% 


7. C, 42.10% 


Na, 14.28-% 


C, 92.3% 


H, 6.43% 


S0 3 24.84% 


H, 7.69% 


0, 51.46% 


Na 2 S0 4 , 44.09% 



INDEX 



(References are to page numbers. The connecting of the names of two chem- 
icals by the word •'and,'' such as "aluminium and nitric acid" indicates an 
experiment in which the reaction between these two substances is studied.) 



Absolute, alcohol. 329 

scale, 192. 239 

temperature, 239 

zero, 238 
Acetic acid. 330 
Acetylene, 315 
Acid, acetic. 330 

citric, 190 

gallic, 190 

hydriodic, 270 

hydrobromic, 267 

hydrochloric. 127 

hydrofluoric. 272 

lactic. 184 

malic, 190 

mucic, 189 

muriatic. 127 

nitric, 170 

oleic. 331 

oxalic. 190 

palmitic, 331 

phosphoric, 314 

prussic, 190 

salts, :;i 1 

silicic, 300 
iric. 331 

sulphuric, 83. 84. 188 

sulphurous, 70, 188 

tartaric, 189 
Acids, tests for. 317 

treating metals with, 281 
Action, catalytic, 53 
Additional problems, 362 

answers 388 



Agent, catalytic, 54 

oxidizing, 310 

reducing, 68 
Air, 39 

and aluminium, 286 

and antimony, 277 

and arsenic, 275 

and bismuth, 295 

and calcium, 158 

and copper, 25 

and iron, 40 

and lead, 279 

and magnesium, 28 

and phosphorus, 32 

and potassium, 150 

and sodium, 130 

and sulphur, 74 

and tin, 285 

and zinc, 27 

composition of, 39, 189, 190, 191 

density of, 243 

fixed, 189 

inflammable, 188 

liquid. 39 

moisture in, 69 

pressure of, 231 

proportion of oxygen in, 34 

specific gravity of, 241, 243 

washing out carbon dioxide 
from, 170 

weight of, 241 
Airslaked lime, 161 
Alchemy, 190 

period of, 184, 185 
Alcohol, absolute, 329 

denatured, 329 



30 1 



392 



INDEX 



(References are to page numbers.) 



Alcohol — Continued 

ordinary, 329 

wood, 329 
Alkali, 132 
Alkaline, 132 
Alkalis, caustic, 189 

mild, 189 
Allotropic forms, 173 
Allotropism, 196 
Allotropy, 73 
Alloy, 297 

fusible, 296, 297, 298 
Alloys, 297 
Alum, 287, 288 

and sodium hydrate, 288 
Alumina, 287 
Aluminium, 286 

and air, 286 

and hydrochloric acid, 287 

and nitric acid, 287 

and sulphuric acid, 287 

bronze, 297, 298 

hydroxide, 288 

oxide, 287 

silicate, 307 

sulphate and sodium silicate, 307 
Alums, 288 
Amalgam, 148 

gold, 149 

lead, 149 

sodium, 147, 148 

zinc, 149 
American Chemical Society, Jour- 
nal of, 254, 378 
Amethyst, 306 
Ammonia, 300, 302, 303 

aqua, 301 

composition of, 298, 301 

fountain, 303 

process, 142 

solubility in water, 303 
Ammonium, 302 

alum, 288 

carbonate and lead nitrate, 284 

chloride and calcium hydrate, 
303 

hydrate, 302 

hydrate and hydrochloric acid, 
302 



Ammonium — Continued 

hydrate and nitric acid, 302 

hydrate and sulphuric acid, 302 

iron alum, 288 

nitrate, heating, 305 

salts, 302 

test for, 303 
Amorphous, 73 
Analysis, 45 

gas, 197 

of unknown salt, 318 

qualitative, 187, 318 

quantitative, 189 

spectrum, 197 
Anhydride, 123 

carbonic, 123 

nitric, 306 

nitrous, 305 

sulphuric, 123 

sulphurous, 123 
Anhydrous, 94 
Animal, charcoal, 116 

kingdom, 115 
Animals, 55 
Anion, 334 
Anode, 332 

Answers to additional problems, 388 
Anthracite coal, 116 
Antimony, 277, 278 

and air, 277 

oxide, 278 

sulphide, 278 

test for, 278 
Apparatus, general, 379 

list of, 1 

individual, 378 
Appendix, 355 
Apples, 190 
Aqua ammonia, 301 
Aqueous vapor, pressure of, 249 
Argentite, 289 
Argon, 39, 335 
Aristotle, 185 
Arrhenius, 333 
Arsenic, 275, 277 

and air, 275 

chloride and hydrogen sul- 
phide, 276 

compounds, 277 



INDEX 



393 



( References are to page numbers.) 



Arsenic — Contin tied 
mirror. 275 
oxide. 275. 277 
oxide and carbon, 276 



poisoninj 



U t 



sulphide, 276 

test for. 276. 318 

white, 276 ■ 
Asbestos, 30 

platinized. 80 
Atomic, period, 184, 191 

theory, 193 

weight of potassium, 258 

weights. 194. 196. 254. 255 

weights, determining. 250 

weights, table of, 255 
Atoms, 193 

number in a molecule, 209 

size of. 194 
Avogadro, 195 

suggestion of, 195, 256 



B 



Babbit's, metal. 297, 298 
Baking soda. 314 
Balance, 188, 190 

gas. 244 

horn-pan, 10 

platform, 8 
Balloons, 189 
Barium, 189, 317 

chloride, per cent of water of 
crystallization in, 316 

compounds. 317 

nitrate. 317 

oxide, 190 

sulphate, 220, 317 

sulphide. 317 
Barometer, 45. 234 

crude, 229 

high-grade, 235 
Barytes, 317 
Base, 317 

Bases, testa for. 317 
Basic, bismuth nitrate, 295 

lead carbonate, 284 

salts, 295 



Bath, steam, 221 
Batteries, Bunsen, 197 

plunge, 59 
Batwing burner, 12 
Beet sugar, 330 
Bell metal, 297, 298 
Bellows, foot, 24 
Bending glass, 15 
Benzine, 330 
Bergman, 189 
Berthollet, 192 

Laws of, 319 
Beryl, 307 
Berzelius, 196 
Bessemer process, 312 
Bismuth, 294 

and air, 295 

and nitric acid, 295 

nitrate, 295 

nitrate, basic, 295 

oxide, 295 

test for, 295 
Bituminous coal, 116 
Bivalent, 204 
Black, 188 

bone, 116 

lead, 115 

magnesia, 190 
Blast, furnace, 311 

lamp, 23 

lamp, lighting, 24 
Bleaching powder and sulphuric 

acid, 275 
Blend, zinc, 28 
Blower, Richards, 24 
Blow- pipe, oxy-hydrogen, 55 
Blue, Prussian, 190 

stone, 104 

vitriol, 104 
Bohemian glass, 307 
Boiling, 20, 22 
Bone black, 116 
Books for reference, 377 
Bottles, catch, 78 
Boyle, 187 

Law of, 187, 229 

period of Robert, 184, 187 

tube, 231 
Brass, 297, 298 



394 



INDEX 



(References are to page numbers.) 



Brimstone, 73 
Brittannia metal, 297, 298 
Brittleness, 20, 21 
Bromides, 267, 269 
Bromine, 266, 269, 274 

test for, 269 

water, 266 

water and potassium iodide, 271 
Bronze, 297, 298 

aluminium, 297, 298 
Bulb tube, 43 
Bunsen, 197 

battery, 197 

burner, 13, 197 

pump, 197 
Burner, bat-wing, 12 

Bunsen, 13, 197 

fish-tail, 12 
Button, silver, 292 



Calces, 191 
Calcite, 167 
Calcium, 158 

and air, 158 

and water, 161 

carbonate, 167 

carbonate and hydrochloric 
acid, 171 

carbonate and sulphuric acid, 172 

carbonate, decomposition of, 168 

carbonate, precipitated, 167 

chloride, 163, 164 

chloride and silver nitrate, 290 

chloride and sodium phos- 
phate, 313 

chloride and sodium silicate, 306 

chloride and sulphuric acid, 164 

fluoride, 272 

fluoride and sulphuric acid, 272 

hydrate, 160 

hydrate and ammonium chlo- 
ride, 303 

hydrate and carbon dioxide, 168 

hydrate and carbonic acid, 167 

hvdrate and hydrochloric acid, 
"162 



Cal c ium — Continued 

hydrate and sulphuric acid, 166 

oxide, 158, 159 

oxide and water, 159 

phosphate, 313 

silicate, 306 

sulphate, 164, 165 

sulphate, solubility of, 165 

test for, 318 
Calorie, 336 
Cane sugar, 330 
Capillary tip, 17 
Carbon, 115 

and arsenic oxide, 276 

and lead oxide, 280 

and oxygen, 116 

gas retort, 116 

monoxide, 121, 123, 189 

oxide, 117 

reduction with, 312 

in organic compounds, show- 
ing, 331 
Carbon dioxide, 117, 123, 189 

and calcium hydrate, 168 

and magnesium, 119 

and potassium, 155 

and sodium hydrate, 142 

and zinc, 120 

in the air, washing out, 170 

preparation on a large scale, 171 

reduction with carbon, 124 

reduction with magnesium, 119 

reduction with potassium, 155 

reduction with zinc, 120 

specific gravity of, 244 

weight of, 244 
Carbonate, test for, 144, 172, 318 
Carbonic acid, 118, 123 

and calcium hydrate, 167 

and potassium hydrate, 154 

and sodium hydrate, 141 

gas, 123 
Carbonic anhydride, 123 
Cast iron, 41, 311 
Catalysis, 53 
Catalytic action, 53 

agent, 54 
Catalyzer, 54 
Catch bottles, 78 



INDEX 



395 



(References are to page numbers.) 



Catching by displacement, 74 
Cathode, 332 
Cation, 334 
Caustic, alkalis. 189 

lime, 159 

lunar, 289 

potash, 151 

soda. 133 
Cavendish, 188. 241 
Chalk, 167 
Changes, chemical. 36 

physical. 36 
Charcoal, 116 

animal, 116 
Charles, law of. 193 
Chemical, changes. 36 

formula from percentage compo- 
sition, 263 

investigation, 174 

method for determining molecu- 
lar weights. 256, 257 
Chemicals, list of. 381 
Chemistry, definition, 183 

Father of, 190 

first great law of. 191. 219 

fourth great law of, 195, 226 

history of, 184 

Industrial, 187, 189 

language of, 198 

Organic, 196, 328 

second great law of, 192, 221 

Technical, 187 

theory of, 183 

third great law of, 193, 222 
China, 308 

Chloride, test for. 318 
Chlorine, 124. 125. 190, 273, 274, 
275 

and hydrogen, 126 

;uid sodium. 136 
Chlorine water, 126. 268 

and potassium bromide, 268 

and potassium iodide, 271 
Choke damp. 12:'. 
Chromium, 285 
Chrome, alum, 288 

yellow, 285 
Cinnabar, 15 
Citric acid. 190 



Clay, 286, 307 

Cleaning deflagrating spoon, 34 

Coal, 116 

anthracite, 116 

bituminous, 116 

gas, 315 

hard, 116 

soft, 116, 314 
Coal-tar industry, 196 
Coefficients, 201 
Color of flame, 13, 90 
Colored glass, 307 
Coin, nickel, 297, 298 

silyer, 297, 298 
Combining number, 194, 248 

of magnesium, 247 
Combustion, 56 

spontaneous, 56 

theory, 190 
Component parts, 77 
Composition, of air, 39, 189, 190, 191 

of ammonia, 298, 301 

of water, 60 
Compound substances, 36 
Concentrated solution, 324 
Concentrating, 324 
Concentration, 324 
Conclusion, 338 
Connector, rubber, 18 
Consistency, 20, 23 
Constant weight, 53 
Conservation of mass, law of, 191, 

219 
Contact process, 85 
Converter, 312 
Cooking soda, 314 
Copperas, 99 
Coral, 168 

Corpuscular theory, 187 
Corundum, 287 
Copper, 19, 27 

aceto-arsenite, 277 

and air, 25 

and nitric acid, 179 

and silver nitrate, 292 

and sulphur, 105 

and sulphuric acid, 101 

arsenite, 277 

nitrate, 182 



396 



INDEX 



(References are to page numbers.) 



Copper — Continued 

oxide, 27 

oxide and hydrogen, 312 

sulphate, 103, 104 

sulphate and hydrogen sul- 
phide, 114 

sulphide, 106 
Cream, of lime, 161 

of tartar, 189 
Crown glass, 307 
Crucible process, 312 
Crude barometer, 229 
Cry of tin, 285 
Cryolite, 273, 286 
Crystalline form, 20, 195 
Crystallization, 91 

evaporating to, 91, 324 

water of, 93 
Curie, 335 

Curve of solubility, 322 
Cut glass, 307 
Cutting glass, 14 
Cyanogen, 190 



Disintegration hypothesis, 335 

Diamond, 115 

Dilute solution, 324 

Displacement, catching by, 74 

Dissociation, 333 

Distillate, 176 

Distillation, 69, 176 

Double substitution, 146 

Drawing out glass, 16 

Dry soda, 142 
tests, 317 

Drying precipitates, 52 

Dryness, evaporating to, 324 

Ductility 20, 21 

Dulong, 195 
and Petit, 252 
and Petit, law of, 195 

Dumas, 197 

Dynamo, 59 



E 



Dalton, 192, 223, 236, 247 

law of, 193, 236 
Davy, 197 
Decantate, 102 
Decantation, 102 
Decrepitation, 283 
Decomposition of calcium carbon- 
ate, 168 
Definite Proportions, by volume, 
law of, 195, 226 

by weight, law of, 192, 221 
Deflagrating spoon, cleaning, 34 
Deliquescence, 132 
Denatured alcohol, 329 
Density of air, 243 
Determining formula of a hy- 
drate, 208 

formula of a salt, 207 

molecular weights, 254 
Developing, 290 
Dewar, 335 



Earthenware, 308 
Effervescence, 94 
Efflorescence, 94 
Ekasilicon, 327 
Electrical, furnace, 187 

machine, 55 
Electrochemistry, 197 
Electrode, 332 
Electrolysis, 57, 197, 333 

apparatus, 57 

of water, 56 
Electrolyte, 333 
Element, 35 
Elements, natural family of, 274 

table of properties of, 356 
Emory, 287 

Endothermic reaction, 336 
Epsom salt, 101 
Equations, 37 

for experiments 1-90, 212 

helps in writing, 202 

twin, 135 

writing, 201, 210, 211 
Equilibrium, ,337 



INDEX 



397 



(References are to page numbers.) 



Equivalent, 248 

weights. 248 
Etching glass. 272 
Ether, 329 
Eudiometer, 60 
Evaporating, to dryness, 324 

to crystallization, 91, 324 
Exothermic reaction, 336 
Experimental work, 219, 265 
Extraction of metals from ores, 312 



F 



Factors, 38 

Families of elements, 274 

Faraday, 197 

law of. 197, 332 
Father, of chemistry, 190 

of pneumatic chemistry, 188 
Feldspar, 286, 307 
Ferric, hydroxide, 311 

iron, 309 

iron, reduction to ferrous, 309 

oxide, 311 
Ferroso-ferric oxide, 311 
Ferrous, hydroxide, 311 

iron, 309 

iron, oxidation to ferric, 310 

oxide, 311 
Film, 290 
Filter paper, 49 
Filtrate, 51 
Filtration, 49. 69 
Firepolishing glass, 15 
First great law of chemistry, 191, 

219 
Fish-tail burner, 12 
Fixed air. 189 
Fixing, 290 

Flame, coloration of potassium, 
153 

coloration of sodium, 153 

yellowed by sodium, 90 
Flask generator, 85 
Flint, 306 

glass, 307 
Flowers of sulphur, 73 
Fluorides, 273 



Fluorine, 272, 273, 274 
Fluorspar, 273 
Flux, 311 

Folding filter paper, 50 
Foot bellows, 24 

Formula, chemical, from percent- 
age composition, 263 

of a hydrate, determining, 208 

of a salt, determining, 207 
Formulae, 199 

graphical, 206 

structural, 206 
Fountain, ammonia, 303 
Fourth great law of chemistry, 

195, 226 
Fundamental substance, 196 
Furnace, blast, 311 

electric, 197 

puddling, 312 
Fusibility, 20, 22 
Fusible alloy, 296, 297, 298 



Galena, 279 
Gallic acid, 190 
Garnet, 286, 307 
Gas, acetylene, 315 

analysis, 197 

balance, 244 

carbonic acid, 123 

coal, 315 

how to test the odor of a, 125 

illuminating, 314 

laughing, 305 

marsh, 315 

olefiant, 315 

Pintsch, 315 

volumes in stoichiometry, 261 

water, 315 
Gas-retort carbon, 116 
Gases, 186 

by heat, law of expansion of, 192 

liquefaction of, 335 

specific gravity of, 188, 242 
Gay-Lussac, 195, 226 
Geber, 185 



398 



INDEX 



(References are to page numbers.) 



General, apparatus, list of, 379 

information, 384 

considerations, 35 
Generator, 85 

flask, 85 

test tube, 127 
German silver, 297, 298 
Germanium, 327 
Gibbs, 338 
Glass, 307 

bending, 15 

Bohemian, 307 

colored, 307 

crown, 307 

cut, 307 

cutting, 14 

drawing out, 16 

etching, 272 

firepolishing, 15 

flint, 307 

lead, 307 

lime, 307 

manipulating, 12 
Glauber, 186 
Glauber's salt, 139, 186 
Glowing splinter test, 44 
Glucose, 330 
Glycerine, 331 
Gold amalgam, 149 
Graduate, reading of, 7 
Grape sugar, 330 
Graphical formula, 206 
Graphite, 115 
Green, Paris, 277 

Scheele's, 277 

vitriol, 84, 99 
Gun metal, 297, 298 
Gypsum, 165 



H 



Halogens, 265, 274 
Hard, coal, 116 

soap, 331 

water, 173 

water, test for, 173 
Hardness, 20, 21 



Hardness — Continued 

of water, 173 
Heat, latent, 189 

of reaction, 336 

specific, 252 
Heating to constant weight, 53 
Heavy, spar, 317 

weighing, 11 
Helium, 335 
Helps, in writing equations, 202 

on the experiments, 386 
Hematite, 41 

High-grade barometer, 235 
History of chemistry, 184 
Hofmann, 196 
Holder, for test tube, 31 
Horn-pan balance, 10 
How to weigh, 9 
Hydrate, determining the formula 

of, 208 
Hydrates, insoluble, 288 
Hydriodic acid, 270 
Hydrobromic acid, 267 
Hydrocarbons, 330 
Hydrochloric acid, 127 

and aluminium, 287 

and ammonium hydrate, 302 

and calcium carbonate, 171 

and calcium hydrate, 162 

and iron, 128 

and lead, 280 

and lead nitrate, 282 

and magnesium, 129 

and manganese dioxide, 273 

and potassium hydrate, 154 

and sodium carbonate, 144 

and sodium hydrate, 139 

and sodium silicate, 306 

and silver, 289 

and tin, 286 

and zinc, 127 
Hydrofluoric acid, 272 
Hydrogen, 60, 188, 189 

and chlorine, 126 

and copper oxide, 312 

and oxygen, 188 

and nitric oxide, 298 

and silver nitrate, 293 

and sulphur, 110 



INDEX 



399 



(References are to page numbers.) 



Hydrogen — Continued 
bromide, 207 

chloride. 120, 127 

fluoride. 272 

from water by magnesium. 65 

iodide, 270 

nascent. 293 

nascent and iron sulphate, 309 

from zinc and sulphuric acid, 
85 

proof that its molecule contains 
two atoms. 209 

reduction by, 313 

showing presence in organic 
compounds. 331 
Hydrogen sulphide, 111, 113 

and arsenic chloride, 276 

and copper sulphate. 114 

and lead nitrate. 308 

and silver nitrate, 291 

and stannic chloride. 308 

and stannous chloride, 308 

and tartar emetic, 278 

and zinc sulphate, 308 
Hydrolysis, 296 
Hydroxyl, 209 
Hypo, 290 



Iron, 40, 41 
alum. 288 
and air, 40 

and hydrochloric acid, 128 
and oxygen. 133 
and sulphur, 108 
and sulphuric acid, 95 
cast, 41, 311 
chloride, 128 
ferric, 309 
ferrous, 309 
hydroxides, 310, 311 
oxidation of ferrous to ferric, 310 
oxide; 41, 311 
pig, 311 

reduction of ferric to ferrous, 309 
rust, 311 
sulphate, 98, 99 

sulphate and nascent hydro- 
gen, 309 
sulphate and nitric acid, 310 
sulphide, 109 

sulphide and sulphuric acid, 112 
weights, 9 
wrought, 41, 311 



J 



Ice, 69 

Iceland spar, 167 

Illuminating gas, 314 

In the cold, 281 

In the hot, 281 

Indestructibility of Matter, law 

of. 191, 219 
Industrial chemistry, 187, 189 
Inflammable air. 188 
Inflammability, 20, 22 
Information, general, 384 
Individual apparatus, list of, 378 
Insoluble hydrates, 288 
Introduction, 1 
Investigation, chemical. 174 
Iodides, 271 
Iodine, 270. 271. 274 
Ions, 333 



Jasper, 306 

Journal of American Chemical 
Society, 254, 378 



K 



Kaolinite, 307 
Kerosene, 330 
Key to the valence, 205 
Kindling temperature, 56 
Kingdom, animal, 115 

mineral, 115 

vegetable, 115 
Krypton, 335 



Lactic acid, L89 
Lamp, blast , 23 



400 



INDEX 



(References are to page numbers.) 



Lampblack, 116 
Language of chemistry, 198 
Latent heat, 189 
Laughing gas, 305 
Lavoisier, 190, 220 
Law, of Boyle, 187, 229 

of Charles, 193 

of chemistry, first great law, 

191, 219 

of chemistry, fourth great law, 

195, 226 
of chemistry, second great law, 

192, 221 

of chemistry, third great law, 

193, 222 

of conservation of mass, 191, 
219 

of Dalton, 193, 236 

of definite proportions by vol- 
ume, 195, 226 

of definite proportions by 
weight, 192, 221 

of Dulong and Petit, 195 

of expansion of gases by heat, 
192 

of Faraday, 197, 332 

of indestructibility of matter 
191, 219 

of mass action, 338 

of multiple proportions, 193, 
222 

Periodic, 197, 325 

verifying a, 219 
Laws, of Berthollet, 319 

of chemistry, 183 

of Raoult, 335 
Lead, 279 

amalgam, 149 

and air, 279 

and hydrochloric acid, 280 

and nitric acid, 280 

and sulphuric acid, 280 

black, 115 

carbonate, 284 

carbonate, basic, 284 

chamber process, 84 

chloride, 282 

chromate, 285 

glass, 307 



Lead — Continued 

monoxide, 280 

oxide and carbon, 280 

oxides, 280 

red, 280 

sulphate, 283 

test for, 309; 318 

tetroxide, 280 

tree, 284 

white, 284 
Lead nitrate, 281 

and ammonium carbonate, 284 

and hydrochloric acid, 282 

and hydrogen sulphide, 308 

and potassium chromate, 284 

and sulphuric acid, 283 

and zinc, 283 

heating, 283 
LeBlanc process, 142, 155 
Lemons, 190 
Libavius, 186 
Liebig, 196 
Light weighing, 11 
Lighting, blast lamp, 24 

hydrogen with safety tube, 88 
Lignite, 116 
Lime, 159, 161 

air-slaked, 161 

caustic, 159 

cream of, 161 

glass, 307 

kilns, 159 

light, 159 

milk of, 161 

phosphate, 38 

slaked, 160 

unslaked, 160 
Limestone, 167 
Lime-water, 161 

and sulphuric acid, 166 

preparation of, 162 
Limonite, 41 
Liquid, air, 39 

chlorine, 197 
Liquefaction of gases, 335 
List, of apparatus, 1, 378, 379 

of properties, 20 
Litharge, 280 
Litmus, 77 



INDEX 



401 



(References are to page numbers.) 



Litmus — Contin ued 

paper. 77 
Lubricating oil. 330 
Lunar caustic. 2S9 
Luster, 20. 21 
Lye, 133, 152 

potash. 152 

soda, 133 



M 



Magnesia, black, 190 
Magnesium, 28, 30 

and air, 28 

and carbon dioxide, 119 

and hydrochloric acid, 129 

and nitric acid, 178 

and oxygen, 54 

and sulphuric acid, 99 

and water. 65 

chloride. 129 

combining number of, 247 

nitrate. 179 

oxide, 29 

sulphate, 100, 101 
Magnetite, 41 
Magnifying glass, 71 
Malic acid. 190 
Malleability, 20, 21 
Manganese, 190 

dioxide, and hydrochloric acid, 
273 

oxide, 190 
Manipulating glass, 12 
Marble, 107 
Marsh gas, 315 
Mass, action, law of. 338 

law of conservation of, 191, 
219 
Massicot, 280 
Matter, law of indestructibility 

of. 191, 219 
Measuring, 4 

rlf'-k. 5 

glass vessels, 6 
Medical period, 184, 180 
Melting point of phosphorus, 31 



Mendeleeff, 197, 325, 327 
Mercury, 41, 45 

and sodium, 147 

and sulphur, 107 

oxide, 43 

oxide, heating, 43 

sulphide, 107 
Metal, Babbit's, 297, 298 

bell, 297, 298 * 

Brittannia, 297, 298 

gun, 297, 298 

speculum, 297, 298 

type, 297, 298 

valve, 297, 298 

white, 297, 298 
Metallic sulphide, by precipita- 
tion, 113, 308 
Metals, extraction from ores, 312 

transmutation of, 185 

treating with acids, 281 
Metathesis, 146 
Metathetical change, 146 
Meteorites, 41 
Metric system, 4 
Meyer, 197, 325 
Mica, 286, 307 
Mild alkalis, 189 
Milk, 189 

of lime, 161 

sugar, 190 
Mineral kingdom, 115 
Minium, 280 
Mirror, arsenic, 275 
Mitscherlich, 195 
Modern period, 184, 332 
Modifications of sulphur, 70 
Moissan, 197 
Moisture in the air, 69 
Molecule, definition, 193 

number of atoms in, 209 

size of, 194 
Molecular weight, potassium chlo- 
rate, 258 

potassium chloride, 258 
Molecular weights, by the chemical 
method, 257 

by the physical method, 256 

determining, 254 
Mortar, 161 



402 



INDEX 



(References are to page numbers.) 



Mucic acid, 189 

Muck, 116 

Multiple proportions by weight, 

law of, 193, 222 
Muriatic acid, 127 



N 



Naphtha, 330 
Nascent, hydrogen, 293 

state, 293 
Natural family of elements, 274 
Negative, 291 
Neon, 335 
Neutral, 138 
Neutralization, 138 
Newlands, 325 
Nickel coin, 297, 298 
Niter, 175 

and sulphuric acid, 175 
Nitrates, heating, 283 
Nitric, anhydride, 306 

oxide, 180, 181 

oxide and hydrogen, 298 
Nitric acid, 176, 181 

a good oxidizing agent, 310 

and aluminium, 287 

and ammonium hydrate, 302 

and bismuth, 295 

and copper, 179 

and iron sulphate, 310 

and lead, 280 

and magnesium, 178 

and potassium hydrate, 182 

and silver, 289 

and tin, 286 
Nitrogen, 39, 189 

dioxide, 181, 189, 305 

monoxide, 180, 181, 305 

oxides, 305 

peroxide, 181, 305 
Nitrous, anhydride, 305 

oxide, 305 
Non-electrolyte, 333 
Normal salts, 295 
Notebook, 1 
Nutgalls, 190 



Octahedral sulphur, 70, 73 
Odor, 20, 22 

of gas, how to test, 125 
Oil, lubricating, 330 

of vitriol, 84 
Oils, 331 
Olefiant gas, 315 
Oleic acid, 331 
Onyx, 168 
Opacity, 20, 21 
Opal, 306 

Open hearth process, 312 . 
Ordinary alcohol, 329 
Ores, 312 

definition, 27 

'extraction of metals from, 312 
Organic chemistry, 196, 328 

compounds, showing carbon in, 
331 

compounds, showing hydrogen in, 
331 
Osmotic pressure, 336 
Ostwald, 335 

Outline of year's work, 339 
Oxalic acid, 190 
Oxidation, 27, 55 

of ferrous iron to ferric, 310 

slow, 56 

theory, 190 
Oxide, 55 

Oxidized silver, 292 
Oxidizing agent, 310 
Oxygen, 26, 39, 55, 189, 190 

and carbon, 116 

and hydrogen, 188 

and magnesium, 54 

and phosphorus, 55 

and sulphur, 74 

and sulphur dioxide, 78 

and zinc, 54 

preparation from mercury ox- 
ide, 43 

preparation from potassium 
chlorate, 45, 46 

proportion in air, 34 

testing for 44 



INDEX 



403 



(References are to page numbers.) 



Oxyhydrogen blowpipe, 55 
Ozone, 55 



Palmitic acid, 331 
Paper, filter. 49 

litmus. It 

turmeric, 133 
Paracelsus, 186 
Paraffine, 330 
Paris, green. 277 

plaster, 165 
Parts, component, 77 
Pearlash, 155 
Peet, 116 
Pentavalent, 204 

Percentage composition from for- 
mula. 2(34 
Period, Modern. 332 

of alchemy. 184. 185 

of Robert* Boyle, 184, 187 
Periods in the history of chemis- 
try. 184 
Periodic, law, 197, 325 

system. 325 

table, 197, 326 
Permanently hard water, 173 
Petit, 195 
Petroleum, 330 
Pewter, 297, 298 
Pfeffer, 336 
Phase rule, 338 
Philosophers, 185 

stone, 185 

wool, 28 
Phlogiston, 187 

period, 184, 187 

theory, 190 
Phosphate of lime, 38 
Phosphates, 314 
Phosphoric acid, 314 
Phosphorus, 30, 38 

and air. 32 

and oxygen, 55 

melting point, 31 

oxide 34 



Photography, 290 
Physical, changes, 36 

method for determining molecu- 
lar weights, 256 
Pig iron, 311 
Pintsch gas, 315 
Plants, 55 
Plaster, 161 

of Paris, 165 
Plastic sulphur, 72 
Plate, 290 

Platform balance, 8 
Platinized asbestos, 80 
Platinum sponge, 80 

tip for flames, 90 
Plumbago, 115 
Plunge batteries, 59 
Pneumatic chemistry, Father of, 188 

period, 184, 188 

trough, 189 
Poison, 42 
Porcelain, 308 
Positive, 291 
Potash, 151 

alum, 288 

caustic, 151 

chrome alum, 288 

lye, 152 
Potassium, 149, 150 

and air, 150 

and carbon dioxide, 155 

and water, 152 

atomic weight, 258 

carbonate, 155 

chlorate, 45 

chlorate, heating, 45, 46 

chlorate, molecular weight of, 258 

chloride, 154 

chloride, molecular weight, 258 

chromate and lead nitrate, 284 

flame coloration, 153 

nitrate, 183 

oxide, 150 

oxide and water, 150 

sulphate, 153, 154 

test for, 318 
Potassium bromide, 268 

and chlorine water, 268 

and silver nitrate, 290 



404 



INDEX 



(References are to page numbers.) 



Potassium bromide — Continued 

and sulphuric acid, 269 
Potassium hydroxide, 151 

and carbonic acid, 154 

and hydrochloric acid, 154 

and nitric acid, 182 

and sulphuric acid, 152 
Potassium iodide, 270 

and bromine water, 271 

and chlorine water, 271 

and silver nitrate, 290 
Pottery, 307 
Pouring down a rod, 51 
Practical questions, 374 
Precipitate, 51, 164 
Precipitated calcium carbonate, 

167 
Precipitates, drying, 52 

washing, 52 
Precipitation of metallic sulphides, 

113, 308 
Preliminary work, 3 
Pressure, of air, 231 

of aqueous vapor, 249 

osmotic, 336 

standard, 244 
Priestley, 189 
Printing, 291 
Prismatic sulphur, 72 
Problems, 225, 226, 234, 240, 241, 
261, 263, 264, 265 

additional, 362 

answers to additional, 388 
Process, ammonia, 142 

Bessemer, 312 

contact, 85 

crucible, 312 

lead chamber, 84 

LeBlanc, 142, 155 

open hearth, 312 

Siemans-Martin, 312 

Solway, 142, 155 

Thomas-Gilchrist, 312 
Products, 38 

Properties, of elements, table of, 
356 

how to test for, 20 

list of, 20 
Property, definition, 19 



Proust, 192, 222 
Prout, 196 
Prussian blue, 190 
Prussic acid, 190 
Puddling furnace, 312 
Purification, by boiling, 69 

by distillation, 69 

by filtration, 69 

of water, 69 
Purifiers, 315 
Pyrites, 41 

Q 

Qualitative analysis, 187, 192, 318 

Quantitative analysis, 189 

Quanti valence, 203 

Quartz, 306 

Questions, practical, 374 

Quicklime, 159 

Quicksilver, 45 



R 



Radium, 335 
Ramsay, 39, 335 
Raoult, 335 

laws of, 335 
Rayleigh, 39 
Reaction, endothermic, 336 

exothermic, 336 

heat of, 336 

reversible, 337 
Reading graduate, 7 
Recrystallization, 103 
Red, lead, 280 

Venetian, 311 
Reducing, agent, 68 

carbon dioxide with magne- 
sium, 119 

carbon dioxide with potassium, 
155 

carbon dioxide with zinc, 120 
Reduction, 68 

by carbon, 312 

by hydrogen, 313 

of ferric iron to ferrous, 309 



INDEX 



405 



(References are to page numbers.) 



Reference hooks, 377 
Replacing numbers, 248 
Reversible reactions, 337 
Rhombic sulphur, 70, 73 
Rhumkorff coil, 61, 65 
Richards blower, 24 
Richter, 191 
Rock, crystal, 306 

salt, 140 
Roll sulphur, 73 
Rubber connector, 18 
Rubbery sulphur, 72 
Ruby, 287 

Rule of stoichiometry, 261 
Rust, 311 
Rutherford, 189, 335 



Safety tube, 88 
Sal soda, 142 
Salt, 140 

analysis of unknown, 318 

definition, 101 

determining the formula of, 207 

Epsom, 101 

Glauber's, 139, 186 

rock, 140 

sea, 140 

solubility in water, 320 

table. 140 
Saltpeter, 175 
Salts, acid, 314 

ammonium, 302 

basic. 295 

normal, 295 
Sand, 306 

Saturated solution, 320, 324 
Saturation, 324 
Sawdust, 190 
Scale, absolute. 239 
Scheele, 189 
Scheele's green, 277 
Sea salt. 140 
Sealing the tube. 88 
Second izt<-.\\ law of chemistry, 

192. 221 
Sexivalent, 204 



Siderite, 41 

Siemans-Martin process, 312 
Silicates, 306 
Silicic acid, 306 
Silicon dioxide, 306 
Silver, 289 

and hydrochloric acid, 289 

and nitric acid, 289 

and sulphuric acid, 289 

bromide, 290 

button, 292 

chloride, 290 

chloride and hydrogen, 293 

coin, 297, 298 

German, 297, 298 

iodide, 290 

oxide, 293 

oxidized, 292 

sulphide, 291 

test for, 318 

tree, 292 
Silver nitrate, 289 

and calcium chloride, 290 

and copper, 292 

and hydrogen sulphide, 291 

and potassium bromide, 290 

and potassium iodide, 290 

and sodium hydrate, 293 
Simple substance, 35 
Slag, 311 
Slaked lime, 160 
Slate, 286 

Slow combustion, 56 
Smaller weights, 10 
Snapping back, 14 
Snow crystals, 70 
Soap, 330 

hard, 331 

soft, 331 
Soapstone, 30 
Soda, 142 

alum, 288 

baking, 314 

caustic, 133 

cooking, 314 

dry, 142 

lye, 133 

sal, 142 

washing, 142 



406 



INDEX 



(References are to page numbers.) 



Soda — Continued 

yellows flame, 90 
Soddy, 335 
Sodium, 129, 130 

amalgam, 147, 148 

and air, 130 

and chlorine, 136 

and mercury, 147 

and water, 134 

bicarbonate, 314 

carbonate, 141, 142 

carbonate, acid, 314 

carbonate and hydrochloric 
acid, 144 

carbonate and sulphuric acid, 
143 

chloride, 137, 140 

chloride and sulphuric acid, 145 

compounds, 149 

dioxide, 131 

flame coloration, 153 

hyposulphite, 290 

monoxide, 131 

oxide, 131 

oxide and water, 131 

peroxide, 131 

phosphate and calcium chlo- 
ride, 313 

silicate and aluminium sulphate, 
307 

silicate and calcium chloride, 
306 

silicate and hydrochloric acid, 
306 

stearate, 331 

sulphate, 138, 139 

test for, 318 
Sodium hydroxide, 132, 133 

and alum, 288 

and carbon dioxide, 142 

and carbonic acid, 141 

and hydrochloric acid, 139 

and silver nitrate, 293 

and sulphuric acid, 137 
Soft, coal, 116, 314 

soap, 331 

water, 173 
Solder, 297, 298 
Solubility, 20, 22, 324 



Solub ility — Continued 

curve of, 322 

of a salt in water, 320 

of ammonia in water, 303 

of calcium sulphate, 165 
Solute, 324 
Solution, 324 

concentrated, 324 

dilute, 324 

in water, 69 

saturated, 324 

supersaturated, 325 

theory of, 333 
Solvay process, 142, 155 
Solvent, 324, 333 
Soot, 116 
Spar, heavy, 317 

Iceland, 167 
Specific gravity, 22, 243 

of air, 241, 243 

of carbon dioxide, 244 

of gases, 188, 197, 242 

standard for, 69 
Specific heat, 252 
Spectrum analysis, 197 
Speculum metal, 297, 298 
Sponge, platinum, 80 
Spontaneous combustion, 56 
Stahl, 187 
Stalactites, 168 
Stalagmites, 168 
Standard, conditions, 244 

pressure, 244 

temperature, 244 

Stannic, chloride and hydrogen 
sulphide, 308 

tin, 308 

Stannous, chloride and hydrogen 
sulphide, 308 

tin, 308 
Starch, 330 
Stas, 196 
State, 20 

nascent, 293 
Steam bath, 221 
Stearic acid, 331 
Steel, 41, 311 
Stibnite, 278 
Stoichiometry, 259 



1XDEX 



407 



(References are to page numbers.) 



Stoichiometry — Contin wed 
involving gas volumes, 261 
rule of, 261 

Stone, blue. 104 
Philosopher's, 185 

Stone-ware, 308 
S. T. P., 244 
Structural formulae, 206 
Substance, elementary, 35 

compound. 36 

fundamental, 196 

simple, 35 
Substitution, 99 

double, 146 
Sugar, beet, 330 

cane, 330 

grape, 330 

milk, 190 
Suggestion of Avogadro, 195, 

256 
Suggestions to the teacher, 382 
Sulphate, test for, 318 
Sulphates in general, 95 
Sulphides, precipitation of, 113, 

308 
Sulphur, 70. 73 

and air, 74. 

and copper. 105 

and hydrogen, 110 

and iron. 108 

and mercury, 107 

and oxygen, 74 

and zinc, 107 

dioxide. 75. 7 V . 

dioxide and oxygen, 78 

flowers. 73 

modifications of, 70 

octahedral, 70. 73 

plastic, 73 

prismatic, 72 

rhombic. 70. 73 

roll. 73 

rubbery. 72 

trioxide, 82 

viscous, 71 
Sulphuretted hydrogen, 111 
Sulphuric acid, 83, 84, 188 

and aluminium, _' v 7 

and ammonium hydrate, 302 



Sulphuric acid — Continued 
and bleaching powder, 275 
and calcium carbonate, 172 

and calcium chloride, 164 

and calcium fluoride, 272 

and calcium hydrate, 166 

and copper, 101 

and iron, 95 

and iron sulphide, 112 

and lead, 280 

and lead nitrate. 283 

and lime-water, 166 

and magnesium, 99 

and niter, 175 

and potassium bromide, 269 

and potassium hydrate, 152 

and silver, 289 

and sodium carbonate, 143 

and sodium chloride, 145 

and sodium hydrate, 137 

and tin, 286 

and zinc, 85 

and zinc oxide, 104 

and zinc sulphide, 114 

on oxides, 105 

on sulphides, 115 
Sulphuric anhydride, 123 
Sulphurous acid, 76, 78, 188 

anhydride, 123 
Supersaturated solution, 325 
Surface tension, 42 
Symbols, 198, 200 

Berzelius', 196 

Dalton's, 194 
Synthesis, 109 



Table, of properties of elements, 356 

Periodic. 197, 326 

salt, 140 
Talc, 30 
Tar, 116 

Tare weighing, 12 
Tartar, cream of, 189 

emel ic, 278 

emetic and aydrogen sulphide, 
27S 



408 



INDEX 



(References are to page numbers.) 



Tartaric acid, 189 

Taste, 20, 22 

Teacher, suggestions for, 382 

Technical chemistry, 187 

Temperature, absolute, 239 

kindling, 56 

standard, 244 
Temporarily hard water, 173 
Tenacity, 20, 21 
Tension, of aqueous vapor, 193 

surface, 42 
Test, for ammonium, 303 

for antimony, 278 
. for arsenic, 276, 318 

for bismuth, 295 

for bromine, 269 

for calcium, 318 

for carbonate, 144, 172, 318 

for chloride, 318 

for lead, 309, 318 

for potassium, 318 

for silver, 318 

for sodium, 318 

for sulphate, 318. 

for zinc, 309 

paper, how to use, 138 

tube, generator, 127 

tube, holder, 31 
Testing, for hard water, 173 

for hydrogen, 88 

for oxygen, 44 

for properties, 20 

odor of a gas, 125 
Tests, dry, 317 

for acids, 317 

for bases, 317 

for tin, 309 

wet, 317 
Tetravalent, 204 
Text-books for reference, 377 
Theories are not facts, 184 
Theory, atomic, 193 

combustion, 190 

corpuscular, 187 

electrochemical, 196 

of chemistry, 183 

of solution, 333 

oxidation, 190 
Thermochemistry, 336 



Thermometers, 45 

Third great law of chemistry, 193, 222 

Thomas-Gilchrist process, 312 

Thunder storms, 55 

Tin, 285 

and air, 285 

and hydrochloric acid, 286 

and nitric acid, 286 

and sulphuric acid, 286 

cry, 285 

stannous and stannic, 303 

tests for, 309 
Tinstone, 285 
Toning, 295 

Torricelli's vacuum, 230 
Tourmaline, 307 
Transferring a dry powder to a 

flask, 48 
Translucency, 20, 21 
Trausmutation of metals. 185 
Transparency, 20, 21 
Tree, lead, 284 

silver, 292 
Trituration, 107 
Trivalent, 204 
Trough, pneumatic, 189 
Tube, Boyle, 231 

bulb, 43 

safety, 88 

sealing the, 88 
Turmeric, 134 

paper, 133 
Twin equations, 135 
Type metal, 297, 298 



U 



Univalent, 204 
Unslaked lime, 160 
Using test papers, 138 



Vacuum, Torricelli's, 230 
Valence, 203 
key to, 205 



INDEX 



409 



(References are to page numbers.) 



Valency, 203 
Valentine, 185 
Valve metal. 297. 298 
Van Helmont, 186 
Van't Hoff, 336 
Vapor tension, 193 
Vapors, 1S6 
Vaseline, 330 
Vegetable kingdom, 115 
Venetian red. 311 
Verifying a law, 219 
Vernier, 235 
Viscous sulphur, 71 
Vitriol, blue, 104 

green, 84. 99 

oil of, 84 

white, 95 
Volatility, 20. 23 

Volume, law of definite propor- 
tions by. 195, 226 
Volumes of gases in stoichiometry, 
261 



W 



Wash-bottle, 17 

Washing, out carbon dioxide in 
air. 170 

precipitates, 52 

soda. 142 
Water, 69, 188 

and calcium, 161 

and calcium oxide, 159 

and magnesium, 65 

and potassium, 152 

and potassium oxide, 150 

and sodium, 134 

and sodium oxide. 131 

bromine. 266 

chlorine, 126. 268 

decomposition by electricity, 
56 

gas, 315 

hard. 173 

of crystallization, 93 

of crystallization, per cent of, 

316 
permanently hard, 173 



Water — Continued 

purification of. 69 

reduction by magnesium, 65 

soft, 173 

solubility of a salt in, 320 

solubility of ammonia in, 303 

standard for specific gravity, 69 

temporarily hard, 173 

test for hardness, 173 

volumetric composition of, 60 
Weigh, how to, 9 
Weighing, 8 

heavy, 11 

light, 11 

tare, 12 
Weight, equivalent, 248 

heating to constant, 53 

law of definite proportions by, 

192, 221 

law of multiple proportions by, 

193, 222 
of air, 241 

of carbon dioxide, 244 

of potassium, atomic, 258 

of potassium chlorate, 258 

of potassium chloride, 258 
Weights, atomic, 194, 196, 254 

determining atomic, 250 

determining molecular, 254 

iron, 9 

molecular, by chemical method, 
257 

molecular, by physical method, 
256 

smaller, 10 

table of atomic, 255 
Wet tests, 317 
White, arsenic, 276 

lead, 284 

metal, 297, 298 

vitriol, 95 
White-wash, 161 
Winkler, 327 
Witherite, 317 
Woehler, 196, 328 
Wood alcohol, 329 
Wool, philosopher's, 28 
Work, experimental, 219, 265 

outline of year's, 339 



410 



INDEX 



(References are to page numbers.) 



Writing equation, 201, 210, 211 

helps in, 202 
Wrought iron, 41, 311 



Xenon, 335 



Yellow, chrome, 285 



Zinc, 27, 28 
amalgam, 149 



Z inc — Continued 
and air, 27 

and carbon dioxide, 120 
and hydrochloric acid, 127 
and lead nitrate, 283 
and oxygen, 54 
and sulphur, 107 
and sulphuric acid, 85 
blend, 28 
chloride, 128 
oxide, 28 

oxide and sulphuric acid, 104 
sulphate, 92, 95 
sulphate and hydrogen sulphide, 

308 
sulphide, 107 

sulphide and sulphuric acid, 114 
test for, 309 



BY THE SAME AUTHOR 

TABLES OF PROPERTIES 

OF OVER FIFTEEN HUNDRED COMMON 
INORGANIC SUBSTANCES 



THE properties given include state, color, luster, crys- 
talline form, deliqescence, efflorescence, stability in 
air, action on test paper, melting point, behavior 
when heated, solubility in water, alcohol and acids, and 
any other properties that are characteristic of the substance 
in hand. For convenience, the formulae, the chemical 
names and the common names are all given. 

The metals are arranged by groups as usually studied 
in Qualitative Analysis. The salts are arranged alpha- 
betically under each metal. 

Non-metals and rare metals are in a section by them- 
selves. Acids also are in a separate section coupled with 
condensed statements relating to the salts of each acid. 

An invaluable reference book, adapted both for rapid 
reference and for comparative study. 

8 vo. x -f 144 pages. Cloth $3.00 (prepaid) 
Specimen pages free 



EXETER BOOK PUBLISHING COMPANY 
Exeter, N. H. 



OCT 6 1909 



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